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HIGH SCHOOL 
CHEMISTRY 



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AMERICAN SOHOOL 

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Book *_,_ 



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COFffilGHT DEPOSIT. 



A STUDY COURSE 

IN 

ELEMENTARY CHEMISTRY 



Prepared Especially for the Instruction 

and Training of Students of the 

American School 



EUGENE E. GILL, Ph.D. 

ASSOCIATE PROFESSOR OF GENERAL CHEMISTRY 
ARMOUR INSTITUTE OF TECHNOLOGY 



B. B. FREUD, B.S., Ch.E. 

ASSOCIATE PROFESSOR OF ORGANIC CHEMISTRY 
ARMOUR INSTITUTE OF TECHNOLOGY 



AMERICAN SCHOOL 

CHICAGO U.S.A. 






COPYRIGHT, 1921, BT 

AMERICAN SCHOOL 



COPYRIGHTED IN GREAT BRITAIN 
ALL RIGHTS RESERVED 



SEP2871 

§>CU6245 43 



A STUDY COURSE 

IN 

ELEMENTARY CHEMISTRY 



DIRECTIONS TO STUDENTS 

Scope of book. This book is not to be taken as a complete 
text in elementary chemistry. It is frankly compiled for use with 
a textbook. In offering it, it is hoped that the lack of labora- 
tory work may be compensated for by the frequent citation of chem- 
ical reactions which have an application in daily life and industry. 
The purpose has been to make this book, in so far as it is possible, 
take the place of class instruction. Little new subject matter has 
been introduced. Much repetition will be found. This was felt 
to be necessary to give proper emphasis to certain topics. 

The course does not claim to be the full equivalent of the 
usual high school course in elementary chemistry where laboratory 
work accompanies the study of the textbook; it does, however, 
cover essentially the same topics. It is believed that a study of 
these pages in connection with the reading of the reference text- 
book will give the student such general knowledge of the principles of 
chemistry as will be useful in forming a part of his general education. 

Method of using book. "An Elementary Study of Chemistry,' , 
second revised edition, by McPherson and Henderson, has been 
selected as the reference textbook; all references to pages occurring 
in this book refer to pages in McPherson and Henderson. Each 
of the twenty-five lessons assigns certain chapters in the textbook 
for study. These chapters should be read first, then the lesson in 
this study course studied, at the same time rereading those parts 
of the textbook that are referred to by pages. After careful study 
of the lesson and the corresponding chapters in McPherson and 
Henderson, the student should prepare the Exercises at the end of 
the lesson in this book, following the directions given in the note 
preceding the Exercises for Lesson I. The Exercises should be sub- 
mitted for grading and correction after each five lessons. 



2 ELEMENTARY CHEMISTRY 

LESSON I 

MATTER AND ENERGY; VARIETIES OF MATTER 
MATTER AND ENERGY 

Assignment: "An Elementary Study of Chemistry," McPherson 
and Henderson, Chapter I 

Science. Chemistry is a study of substances and their 
behavior toward other substances. It deals with the things with 
which we come in daily contact and thus has an important part 
in the understanding of the changes that take place in our daily 
lives. It is one of the natural sciences. Science may be defined 
as a study of the facts of nature, of the laws that state these facts, 
and the theories that explain the laws. 

When men first began to observe the phenomena of nature as 
it existed and changed about them, only the one general subject 
of natural science was necessary to be considered, but soon the 
facts to be studied became so numerous that a division into 
branches of science was necessary. Today, to name only the 
most important sciences, we have botany, anatomy, geology, 
astronomy, chemistry, and physics. 

Physics and chemistry. Of these sciences each one is more or 
less related to every other science. No one can be independent of 
the others, but we find that two, chemistry and physics, are much 
more intimately related than the others. These two are more 
fundamental as regards their subject matter. Botany deals with 
plants, anatomy with the animal body, geology with the earth's 
crust, astronomy with the heavenly bodies, but both chemistry 
and physics are more fundamental and study that which makes up 
these other bodies. 

Matter. There are two fundamental things or ideas which we 
have to consider in studying the facts of nature. One is that to 
which we refer as matter. This composes all the things around 
us — chairs, books, our bodies, the earth, air, water, etc. The term 
matter is not easily defined. Matter is recognized by its proper- 
ties, such as mass, inertia, impenetrability, etc. Of these, mass 
serves best to give a definition and matter may be~said to be any- 
thing that has mass, or occupies space. This mass is usually deter- 



ELEMENTARY CHEMISTRY 3 

mined by weighing the substance. Weight varies with altitude 
and latitude, mass does not. Therefore, weight is not a good 
definition of matter. 

Energy. The other fundamental thing is energy. A little 
observation will show many examples of changes taking place in 
matter. Water freezes, wood burns, iron rusts, milk sours, alcohol 
evaporates, rocks disintegrate, animals and plants die and decay. 
The thing that brings about these changes in matter is called 
energy. Energy is sometimes defined as that which does work. 

Changes in matter. A further study of these changes will 
show that they are of two kinds. If we take two wires, one made 
of platinum, the other made of magnesium, and hold each in a 
flame, we observe a difference in their behavior. Both give off 
light. The platinum retains its form and when removed from the 
flame returns to its original condition. It looks, feels, weighs, and 
measures the same. The magnesium wire, however, changes to a 
white powder and does not return to its original form and appear- 
ance. This white powder can be shown to weigh more than the 
original magnesium. 

Physical changes. In the first case the platinum took back 
its original form when the original conditions returned. Such a 
change is an example of physical change. Other examples of 
physical change are freezing of water, melting of ice, moving of 
objects, dissolving of sugar. In all these cases it is noticed that 
the change is temporary and lasts only as long as the cause (high 
temperature in some cases) operates. The original properties 
return when this cause is removed. We learn to associate proper- 
ties with the composition of matter. We may then say that a 
physical change is one that does not change the composition of matter. 
Physics is the science that studies physical changes. 

Chemical changes. When the magnesium wire was heated, it 
underwent a permanent change. When the cause (heat) was 
removed, the matter did not return to its original form. A new 
substance was formed with a new set of properties. The metallic 
wire became a white powder which weighed more than the wire. 
This change in properties is associated with a change in the com= 
position of the substance. Such a change is a chemical change and 
may be defined as a change in the composition of matter, recognized 



4 ELEMENTARY CHEMISTRY 

by a permanent change in its properties. Chemistry is the study of 
chemical changes. The burning of fuels is the most common 
chemical change. The rusting of iron, souring of milk, explosion 
of gunpowder, digestion of food, fermentation of sugar, and dis- 
integration of rocks are other chemical changes. 

Conservation and transformation of energy. Energy exists in 
several forms, such as heat, electricity, mechanical energy, and 
light. These forms of energy can be converted into each other. 
The burning of coal gives heat, and this produces steam to run a 
steam engine, which may run a dynamo to give electricity, which 
in a light bulb gives light and heat (see illustration of energy 
transformation, page 7, McPherson and Henderson). 

The question arises, in what form did the energy exist that 
gave the heat when the coal was burned? In the act of burning, 
the oxygen of the air and the carbon of the coal both undergo a 
chemical change. In this change heat energy is liberated. We 
say that this energy existed in the coal and the oxygen as chemical 
energy. 

Every chemical change is accompanied by an energy change 
of some kind; energy is either given off or absorbed, usually as 
heat. Experience has taught us that energy cannot be created or 
destroyed but only transformed from one hind to another. This is 
stated as the law of the conservation of energy. A scientific law 
may be defined as a general statement of facts. 

Conservation of matter. In all the changes which matter 
may undergo, it can be shown that there is neither loss nor gain in 
the amount of matter taking part in the change. As measured by 
weighing, it has been shown that the weights of all the products of 
a change are exactly equal to the weights of all the substances 
entering into the change. We may transform matter but cannot 
change its mass. This is known as the law of the conservation of 
matter and may be stated as follows: In any change to which mat- 
ter is subjected, the mass remains the same. 

One sometimes has to think a second time to see the truth of 
this in the case of burning fuels, where the ash left is obviously 
quite small compared to the mass of fuel used. If, however, 
weight of fuel plus weight of oxygen is compared with weight of ash 
plus weight of gaseous products formed, they will be found to be equal. 



ELEMENTARY CHEMISTRY 5 

VARIETIES OF MATTER 
Assignment: Chapter II, McPherson and Henderson 

Classes of matter. From the point of view of its composi- 
tion, matter may be divided into three classes, namely: elements, 
compounds, and mixtures. As illustrative of these classes, the 
experiments with iron and sulfur and with sugar should be 
studied (pages 13 and 14). 

Heating either the iron or the sulfur without contact with the 
air would result in no chemical change. Intimately mixing them 
without heat still leaves them in a condition where all their 
individual properties are recognizable. Heating them together, 
however, produces a new substance with a set of properties of its 
own and differing from the original properties of iron and sulfur. 
The glow observed at the time of the action is due to heat liber- 
ated from the chemical energy existing in the iron and the sulfur. 
The action is exactly similar to the kindling of coal and oxygen, 
where the coal continues to burn after the fire is started. 

Heating sugar, however, decomposes it into two new sub- 
stances, carbon and water. Passing an electric current through 
water shows that it in time can be decomposed into two gases, 
oxygen and hydrogen. 

Elements. The question arises, can these decompositions be 
carried farther? Continued trials convince us that hydrogen, 
oxygen, carbon, iron, and sulfur cannot be decomposed into other 
substances by any means now known. That is not saying that it 
is impossible to decompose them. Water was at one time thought 
to be undecomposable. Even in recent years it has been shown 
that some substances which were thought to be simple were in 
reality composed of two or more constituents. 

Substances like those mentioned above are called elements. 
An element is a substance which has not as yet been decomposed into 
simpler substances. At present the number of known elements is 
between eighty and ninety. 

Compounds. The iron sulfide formed when the iron and the 
sulfur were heated is a compound. Sugar and water, substances 
that can be decomposed into simpler substances, are compounds. 
This class of substances must be distinguished on the one hand 
from elements and on the other from mixtures. 



6 ELEMENTARY CHEMISTRY 

This is not always easy, but the study of large numbers of 
substances shows that those that can be classed as compounds 
have a constant percentage composition. Thus, water is always 
found to contain 11.19 per cent hydrogen and 88.81 per cent 
oxygen. 

A chemical compound is a substance composed of two or more 
elements in constant proportions, which elements have lost their 
original properties, the new substance having a set of properties of its 
own, and it cannot be separated into its elements by mechanical 
means. 

Mixtures. The substance composed of iron and sulfur before 
it was heated is an example of a mixture. Sand and sugar 
shaken together would be a mixture. Ordinary concrete is 
another. The rock called granite is composed of three distinct 
substances which can be recognized by the naked eye: quartz, 
feldspar, and mica. All three are themselves well-known chemical 
compounds. 

In all these instances, the substances can be separated into 
their constituent parts by mechanical means. The magnet sep- 
arates iron from sulfur. Water would dissolve sugar from sand. 
Crushing the granite would make it possible to pick out the three 
kinds of rock. In these cases each constituent can be recognized 
by its individual properties, even while existing in the mixture. 
Further, the percentage amounts of the various constituents may 
vary, without radically changing the substance. 

A mixture is a substance composed of two or more elements or 
compounds, in proportions that may vary, which still retain their 
original properties and can be separated from each other by mechani- 
cal means. 

Chemical reactions. In the formation of compounds from the 
elements or in the decomposition of compounds into elements or 
simpler compounds, there is a considerable change in properties. 
Products are formed that are unlike the original substances. 

Just as important as this, however, is the energy change that 
takes place. In every chemical change there is a transformation 
of energy. Chemical energy held in the substances is transformed 
into some other form, such as heat, and liberated; or heat, or 
other form of energy, is used and converted into chemical energy, 



ELEMENTARY CHEMISTRY 7 

which is held in the substances produced by the chemical change. 
This energy change is a necessary characteristic of every chemical 
change, or chemical reaction, as it is called. There is always a 
change in the composition of matter at the same time that the 
energy change occurs. 

It is not possible to say what causes chemical changes. 
Energy change and matter change are simultaneous and neither 
can be said to be the cause of the other. To say that chemical 
affinity is the cause of chemical change tells us nothing because no 
one knows what causes chemical affinity. The expression simply 
means a chemical attraction, or tendency to combine. 

It is important, however, to keep in mind that conditions 
have considerable influence in determining whether a reaction will 
or will not take place, and if it does take place, they determine its 
velocity. Applying heat or other forms of energy frequently pro- 
motes the speed of a reaction, either of combination or decompo- 
sition. 

EXERCISES* 

Read Carefully: Write out your answers to these questions. Place your 
name and full address at the head of the paper. Any cheap, light paper may be 
used; write on only one side of the paper. Do not crowd your work, but arrange it 
neatly and make it legible. Do not copy the answers; use your own words so that 
we may be sure you understand the subject. 

1. State the meaning of science, matter, energy, scientific law. 

2. Name and define two kinds of changes in matter. 

3. Define chemistry and physics. 

4. Give a number of examples of physical and chemical changes. 

5. Name all the forms of energy. 

6. Describe an experiment showing their transformation. 

7. State the laws of conservation of energy and of matter. 

8. Describe an experiment showing the formation of a compound of 
iron and sulfur. 

9. Define element, compound, and mixture. 

10. Discuss the nature of chemical reactions and conditions influencing them. 

11. Explain free, or native, state. What is a calorie? 

12. Describe the decomposition of water. 

13. Give the experiment of Lavoisier. 

14. How could you show that the substances formed in burning a candle 
weigh more than the candle? What does this show? 

15. Solve problems 15 and 16, page 11, McPherson and Henderson. 

16. Answer questions 1, 3, 4, 5, 9, and 10, pages 10 and 11. 

17. Answer questions 1, 6, 10, 11, 12, and 13, pages 22 and 23. 

* Prepare this set of Exercises and hold it until those for Lessons II, III, IV, and V are 
also prepared and then send all five sets to the School. 



8 ELEMENTARY CHEMISTRY 

LESSON II 

OXYGEN 
Assignment: Chapter III, McPherson and Henderson 

Importance. The study of chemistry is first of all a study of 
substances. Of these it is natural to study the simple ones, that 
is, the elements, before taking up the compounds. In connection 
with the study of each elementary substance, its compounds will 
be considered. Of the elements, oxygen is by far the most 
important and is therefore studied first. Oxygen occurs in greater 
quantities than any other element. It occurs in the air as an 
element. In combination with other elements it is found in 
water, all forms of living matter, and the crust of the earth. 

Oxygen is the active constituent of the atmosphere in which 
we live and is essential to all animal life. It also supports all the 
burning of fuels which supplies heat. All chemical changes that 
take place in nature, as well as the vast majority of those that 
are carried out in our laboratories and industrial plants, take 
place in an atmosphere of oxygen. The possible effect of the 
action of oxygen on the materials involved in these changes must 
be taken into account. So it is well at the beginning of our study 
to learn as much as we can of the properties and chemical 
behavior of oxygen. 

Discovery. It was the discovery of oxygen as a constituent 
of the air which would support life and the burning of fuels more 
vigorously than air itself would that started the development of 
modern chemistry. Priestly and Scheele made the discovery at 
about the same time. It was Priestly who performed the experi- 
ments with mice, finding that a mouse could not live in air from 
which all oxygen had been removed and that in pure oxygen the 
mouse's activities were so vigorous that the animal soon died of 
exhaustion. 

However, it was Lavoisier who recognized the exact nature of 
the substance and who gave it the name "oxygen." It was 
thought by Lavoisier that all acids contained oxygen, being formed 
by the addition of water to the oxides of non-metallic elements. 
Hence, he called the element oxygen, which means acid producer. 
He proved that the burning of metals resulted in an increase of weight 



ELEMENTARY CHEMISTRY 9 

by an amount equal to the loss in weight on the part of the air 
(see page 4.). This experiment established the relation of oxygen 
to that most important chemical change — combustion, or burning — 
as well as successfully combated the older theory of phlogiston. 

Preparation. In the preparation of oxygen we must look to 
its natural sources, air, water, and other compounds containing it. 

1. Obtaining oxygen from air. Air is a mixture of oxygen 
and nitrogen, both of which gases can be liquefied; but the result- 
ing liquids have different boiling points (see "Liquid air," pages 
108 and 141). If air is liquefied, the nitrogen boils off first, as it 
has the lower boiling point. Oxygen will be left. Oxygen obtained 
in this way is not pure, but it is pure enough for many com- 
mercial purposes. A large amount of oxygen for such uses is 
prepared in this way. 

2. Obtaining oxygen from water. Passing an electric current 
through water which has a little sulfuric acid dissolved in it gives 
two gases, oxygen and hydrogen. This action is sometimes called 
the electrolysis of water (pages 16 and 17). The method gives 
quite pure oxygen and is used for its preparation on a large scale 
when purity is desired. By a similar action oxygen gas is liber- 
ated at one of the poles in certain electroplating processes. 

3. Obtaining oxygen from various compounds. Heating cer- 
tain compounds that contain oxygen will cause it to be liberated; 
not all compounds containing oxygen will give it up when heated. 
A few that yield oxygen when heated are mercuric oxide, potas- 
sium chlorate, sodium peroxide, and potassium nitrate. A com- 
mon method for making oxygen on a small scale is to heat potas- 
sium chlorate. This gives oxygen and a compound of potassium 
and chlorine. This action works faster and at a lower tempera- 
ture if some manganese dioxide is added, although just what part 
the manganese dioxide plays is not known, as no change can be 
noted in it when the action is complete, and as much of it remains 
as was used in the beginning. A substance used to change the 
velocity of an action, without apparently taking any part in it, is 
called a catalytic agent. Other examples of catalytic agents will be 
met with later. 

Priestly obtained his oxygen by heating mercuric oxide, a red 
powder containing mercury and oxygen. This method is slow 



10 ELEMENTARY CHEMISTRY 

and expensive. Sodium peroxide when treated with water will 
yield oxygen and a compound known as sodium hydroxide. This 
is a good quick method to make the element. Examples of com- 
pounds which will not liberate oxygen when heated, though they 
contain considerable quantities of it, are copper oxide, calcium 
carbonate (or limestone), iron oxide, sugar, alcohol, sand, and 
carbon dioxide. 

Properties. The physical properties and the chemical proper- 
ties of substances must always be studied. The latter are some- 
times referred to as chemical conduct. Even though we are 
studying chemical changes, it is necessary to know the physical 
properties of the substances we are dealing with. We cannot see a 
chemical change, and it is only possible to know that one has 
taken place by recognizing that we have obtained a substance 
with a new set of physical properties. Six important physical 
properties should be considered for every substance studied. 
They, in general, correspond to the common physical senses. 
They are physical state, taste, color, odor, density, and solubility. 

Oxygen is a tasteless, colorless, odorless gas which can be 
liquefied. It is slightly heavier than air and slightly soluble in 
water. When liquefied, it boils at —182.9° C. at atmospheric 
pressure. 

Chemical conduct. The chemical conduct of oxygen as a 
supporter of burning, or combustion, is apparent. This, however, 
is only a part of its general tendency to combine with other ele- 
ments or with compounds. The ordinary fuels are organic sub- 
stances containing carbon and hydrogen. Oxygen will combine 
with these elements to form carbon dioxide and water, respec- 
tively. Oxygen will combine by direct addition with nearly all 
the elements. Of the eighty some elements known, only thirteen 
do not combine directly with oxygen. Gold, silver, and platinum 
are among these thirteen. By indirect methods oxygen compounds 
can be made with five of this group of thirteen. This fact 
emphasizes the wide range of its activity. 

Its action is usually not very vigorous at ordinary tempera- 
tures. At such temperatures it does combine with a few ele- 
ments like sodium and phosphorus with a fair velocity, but the 
action is made very rapid by raising the temperature. Usually a 



ELEMENTARY CHEMISTRY 11 

rather high temperature is necessary to cause appreciable union of 
oxygen with other substances. In the ignition of a match the 
necessary heat to ignite the phosphorus or phosphorus compound 
used is generated by friction, the oxygen being supplied by some 
compound like potassium chlorate. In old-fashioned gunpowder 
the combustion consists of the carbon and sulfur of the mixture 
combining with oxygen supplied by potassium nitrate. 

Compounds formed by the union of oxygen with another ele- 
ment are called oxides. The union of oxygen with another sub- 
stance is called oxidation. A substance that supplies the oxygen 
is called an oxidizing agent. Oxidation may be of two kinds, slow 
and rapid. The products of the action are the same in both 
cases and the total heat given off is the same in both. 

However, if the oxidation is rapid enough, the heat will be 
intense enough to produce light. This is called combustion. In 
many cases combustion reactions are carried out for the heat they 
generate and not for the products they form. This is the case in 
all burning of fuel. The temperature at which light is given off, 
or combustion commences, is called the kindling temperature. 

Uses of oxygen. Free oxygen is essential to the life of all 
animals except a few very low forms. Water animals obtain 
oxygen from the air that is dissolved in the water. Oxygen is 
used in the decay of waste organic matter, the decaying substance 
being oxidized into harmless gaseous substances. Certain forms of 
bacterial life are necessary to cause this oxidation to take place. 

An example of purification by oxidation is the method some- 
times used for sewage. The sewage is sprayed into the air (see 
page 34), and bacteria and oxygen destroy the harmful organic 
matter. Oxygen is utilized in some industrial processes and in 
laboratory work. Pure oxygen is used in the treatment of certain 
lung diseases where the patient cannot inhale enough air to sup- 
port life. 

Law of definite composition. In all the chemical changes 
that have been discussed in this and the previous lesson, it must 
be remembered that the change has always been between certain 
fixed weights of the various constituents if a compound is formed; 
or, if a compound is decomposed, it is always found to contain 
exactly the same percentages of its elements. The law of definite 



12 ELEMENTARY CHEMISTRY 

composition may be stated as follows: The composition of a chemi- 
cal compound never varies. This is true no matter where the com- 
pound may be found, nor does the manner of its preparation 
make any difference in its percentage composition. 

EXERCISES 

1. Relate the facts connected with the discovery of oxygen. 

2. Describe the method for making pure oxygen on a commercial scale. 

3. If oxygen on a large scale, but not pure, is required, how is it 
obtained? 

4. Name several compounds that will yield oxygen when heated. Name 
several that will not. 

5. What is meant by catalysis? Name a catalytic agent and the 
circumstances of its use. 

6. Give the physical properties of oxygen. 

7. Discuss the chemical conduct of oxygen. 

8. Define oxide, oxidation, combustion, and kindling temperature. 

9. What is meant by heat of oxidation and combustion? Are they 
the same? Why? 

10. Explain spontaneous combustion. Give examples. 

11. Name the uses of oxygen. 

12. State the law of definite composition. Illustrate with examples. 

13. Solve problems 12, 13, 14, 15, and 16, pages 36 and 37. 

14. Answer questions 1, 2, 7, 8, and 10, page 36. 



LESSON III 

HYDROGEN; GAS LAWS 

HYDROGEN 

Assignment: Chapter IV, McPherson and Henderson 

Discovery and occurrence. The study of hydrogen is second 
in importance only to that of oxygen because it, with oxygen, 
forms our most important chemical compound, water. It was 
isolated from water by Cavendish, and for that reason called 
hydrogen, meaning to produce water. Hydrogen does not occur in 
the elementary form in nature unless in rare traces. In the com- 
bined state it is found in water and in that large class of com- 
pounds called organic, which are the chief constituents of the 
bodies of all plants and animals. It is a constituent of sugar, 
starch, petroleum, natural gas, acids, and bases. 



ELEMENTARY CHEMISTRY 13 

Preparation. 1. Electrolysis of water. Hydrogen is prepared 
from water by the use of the electric current. This method was 
studied under oxygen. 

2. Action of certain metals on water. Some metals like 
sodium and potassium will liberate hydrogen from water at ordi- 
nary temperatures, forming at the same time compounds known 
as sodium or potassium hydroxides. These compounds belong to 
the class called bases and in their impure form are, respectively, 
the soda and potash lyes in common household use. In this prep- 
aration of hydrogen much heat is evolved and sometimes this 
will be sufficient to set fire to the hydrogen gas as it is evolved. 
Other metals, like iron and zinc, will liberate hydrogen from 
water only at high temperatures. Under these conditions the 
oxides of the metals are formed. Such metals as copper and 
silver will not set hydrogen free from water under any conditions. 

3. Action of certain metals on acids. This is the usual labora- 
tory method for making hydrogen. The metal used must be of 
such a nature as to be able to replace hydrogen from the acid. 
Iron and zinc are very suitable for this purpose, but copper and 
silver would not do at all. Theoretically, any acid might be 
used, but some acids are so weak that the evolution of hydrogen 
would be too slow to serve the purpose, while, in the case of 
other acids, secondary reactions take place that use up the 
hydrogen as fast as it is formed. For practical purposes, then, 
hydrochloric and sulfuric acids are the common ones used in the 
preparation of hydrogen. In the use of zinc and hydrochloric acid, 
the products of the action are hydrogen and zinc chloride. Iron 
and sulfuric acid give hydrogen and iron sulfate. For commercial 
uses hydrogen is made from water by using the electric current, 
or by using steam and iron. 

4. Water gas as source of hydrogen. There has recently 
come into use a method of obtaining hydrogen for commercial 
purposes from water gas. This gas is made by passing steam 
over hot carbon in the form of coke or anthracite coal and con- 
sists of a mixture of hydrogen and carbon monoxide. Water gas 
is frequently used for fuel and illuminating purposes. Hydrogen can 
be obtained from it by liquefying the carbon monoxide. Hydrogen 
for balloon uses was made in this way during the European War. 



14 ELEMENTARY CHEMISTRY 

Properties and chemical conduct. Hydrogen is the lightest of 
all the elements, being 15.88 times lighter than oxygen. Because 
of its lightness, it is much used to fill balloons. When pure, it is 
an odorless, colorless, and tasteless gas. It is slightly soluble in 
water. It can be liquefied and solidified. A remarkable property 
of hydrogen is its ability to be absorbed, or occluded, by metals — 
in most cases in traces, but gold, platinum, and palladium occlude 
large volumes of it. 

In its chemical conduct it is much less active than oxygen 
With some elements under proper conditions it forms compounds 
called hydrides. It will combine with oxygen, chlorine, nitrogen, 
and sulfur quite readily. With oxygen and chlorine it forms 
explosive mixtures which are dangerous if not handled carefullv. 
In working with hydrogen it is necessary to prevent its mixing 
with air, for dangerous explosions would occur if a flame came ii. 
contact with such a mixture. When oxygen and hydrogen ai 
mixed in proper proportions and heated to 800 degrees, a violent ex- 
plosion takes place. A jet of hydrogen gas will burn quietly in the 
air, forming water; the flame is almost colorless if free from 
impurities. By weight, 1 part of hydrogen combines with 7.94 
parts of oxygen. 

It is to be noted that chemically hydrogen is the opposite of 
oxygen. It does well the things oxygen does poorly and does not 
do the things oxygen does best. Oxygen combines with most of 
the metals, hydrogen does not; hydrogen combines with the non- 
metals, oxygen does so with less facility. For this reason it is 
usually said that hydrogen does not support combustion. In other 
words, it does not combine with the things that we have learned 
will burn in the air, that is, in oxygen. 

Reduction. Hydrogen has so great a tendency to combine 
with oxygen that it will extract it from many of its compounds 
under suitable conditions. If we pass hydrogen over heated cop- 
per oxide, the oxygen will combine with hydrogen to form water 
and leave copper. The making of iron from iron oxide in a blast 
furnace is a similar reaction in which carbon takes the place of 
hydrogen. The oxide of the metal is reduced. Reduction is the 
removing of oxygen from a compound. The substance used to 
bring about this reduction is called a reducing agent 



ELEMENTARY CHEMISTRY 15 

Reduction and oxidation are opposite processes. The two 
processes take place together, one substance being reduced and the 
other oxidized. Later it will be shown that these terms can be 
used to include elements with properties similar to oxygen and 
hydrogen. 

GAS LAWS 

Assignment: Chapter V, McPherson and Henderson 

Volume of gases. We are all familiar with the variation in 
volume of gases under different conditions of pressure and. tem- 
perature. Toy balloons and rubber tires show this. The gas in 
them contracts under outside pressure and expands with increase 
li temperature, sometimes causing blowouts. Quantities of gases 
iare measured more conveniently by volume than by weighing. 
But, since the volume varies with conditions, the measured volume 
must be corrected to a standard set of conditions for pressure and 
temperature according to the laws of expansion and contraction. 

The law of Boyle. The law of Boyle states the rule for vol- 
ume change when the pressure changes and other conditions 
remain unchanged: The volume of a given mass of any gas varies 
inversely as the pressure, the temperature remaining constant. Thus, 
if a gas measures 500 cc. under a certain pressure, it will measure 
1000 cc. if the pressure is halved, or 250 cc. if the pressure is 
doubled. This law may be expressed by a formula by letting Pi 
and P 2 represent the two pressures and V\ and V% the two vol- 
umes. Then 

n:F,::P,:Pi 
which may be written 

l\ = P2 
V 2 P l 

or 

V 1 P 1 =V 2 P 2 

Standard pressure. We must have a unit of measurement 
for pressure as we have a yardstick to measure length or a 
pound to measure weight. For pressure, the pressure of the 
atmosphere at sea level is chosen. The weight of this column 
of air will support a column of water, of equal cross-sectional 
area, 1033.3 cm. in height. Since 1 gm. of water measures 1 cc, a 



16 ELEMENTARY CHEMISTRY 

column of water 1 sq. cm. in cross-sectional area and 1033.3 cm. 
high weighs 1033.3 gms., and a column of air of the same cross- 
sectional area weighs the same. 

The column of water is too long for convenient use so its 
value is expressed in terms of a denser liquid, mercury. Mercury 
is 13.59 times as dense as water. Divide 1033.3 cm. by 13.59 and 
we have 76 cm. as the height of the column of mercury which one 
atmosphere will support. This is usually read as 760 mm., and 
standard pressure is equal to that exerted by a column of mercury 
760 mm. in height. 

Example. Suppose oxygen is measured as 500 cc. at a pressure of 730 mm. 
What will its volume be at 760 mm.? 

According to Boyle's law, Vi'. Y%'.\ P2'. Pi, 500 cc. and 730 mm. are the 
first conditions, and the unknown volume X, and 760 mm., the second condi- 
tions. Substitute these in the formula and we have 

500: X:: 760: 730 

Solving 

760Z = 500X730 

500X730 
X= * =480.26 cc. 
760 

Absolute temperature. A study of the expansion of gases for 
temperature changes shows that there is an increase in volume of 
2T¥ of the volume at 0° C. for every degree centigrade the tem- 
perature is raised. Correspondingly there is a decrease of sts of 
the volume at zero centigrade for every degree it is lowered. 

If the temperature were lowered 273° below zero, the given 
mass of gas would occupy no volume. This is impossible, but all 
known gases become liquids before reaching this point. Theoreti- 
cally —273° C. becomes the temperature of no volume, and this 
temperature is therefore the absolute zero. A thermometer con- 
structed with this point as zero would measure absolute degrees of 
temperature. A comparison of these degrees with the volume of a 
gas would show them to be directly proportional. 

The law of Gay=Lussac. The above facts may be stated as 
follows: The pressure remaining constant, the volume of a given mass 
of any gas will vary directly as the absolute temperature. This is 
called the law of Gay-Lussac. It may be expressed in formulas as 
follows : 



ELEMENTARY CHEMISTRY 17 



or 

V 2 T 2 
or 

Example. Five liters (5000 cc.) of hydrogen are measured at 27° C. 
What will its volume be at 7° C? 

In solving a problem of this nature it is always necessary to change the 
temperatures which are given in centigrade degrees to absolute degrees by 
adding 273 to the temperature given. Using the formula ViiVtiiTiiT* and 
substituting the appropriate values for the letters, we have 

5000: X:: (27+273): (7+273) 
Solving 

300X = 1400000 
X = 4666.67 cc. 

If both temperature and pressure changes are to be corrected 
for, the result in one case is to be used as the given value in the 
second correction, and each part of the problem solved as shown 
above. 

Example. What will a given quantity of a gas measure at 37° C. and 
700 mm. pressure, if it measures 500 cc. at standard conditions? Standard 
conditions are 0° C. and 760 mm. pressure. 

We may correct for pressure first, applying the formula of Boyle's law, 

Fi:7,::P.:Pi. 

Substituting we have 

500: X:: 700: 760 

Then 

_ 500X760 
%= — ^7 — = 542.8 cc. 
700 

This value now becomes the given volume to be corrected for temperature 

according to the law of Gay-Lussac, which is represented by the following 

formula: ViiVtiiT^T* 

Substituting we have 

542.8: X:: 273: (273+37) 

„ 542.8X310 

X = ^ =616.3 cc. 

273 

Therefore 616.3 cc. is the volume occupied by the gas at 37° C. and 700 mm. 
pressure. 

The Kinetic theory. The gas laws just studied are experi- 
mental facts. But why do all gases behave in such a regular 
manner? An attempt to explain the facts as stated by a law is 
called a theory. These laws are explained by what is known as 



18 ELEMENTARY CHEMISTRY 

the kinetic theory. The word "kinetic" means motion. It is 
supposed that all gases are composed of many small moving par- 
ticles. These are called molecules. The molecules are far apart, 
hence the compressibility of gases. They are moving in straight lines 
with great velocities, striking each other and the sides of the contain- 
ing vessel; this accounts for the pressure the gas exerts and its tend- 
ency to expand. Heat increases the velocity and therefore causes the 
expansion. 

Assuming equal weight for all the molecules of the same gas 
but different weights for molecules of different gases, the kinetic 
theory suggests that equal volumes of all gases, the temperature and 
pressure remaining the same, contain the same number of molecules 
since they exert equal pressure. This statement is known as 
Avogadro's hypothesis and will be referred to in Lesson XVI. 

EXERCISES 

1. Describe three methods for making hydrogen. 

2. State the physical properties of hydrogen. 

3. Discuss the chemical conduct of hydrogen. 

4. Define reduction and give two examples. 

5. Explain why hydrogen sometimes explodes. 

6. What are hydrides? Name four. 

7. Solve problems 3, 12, 13, 14, and 15, page 51. 

8. State the law of Boyle; of Gay-Lussac. 

9. Discuss the selection of a standard pressure. 

10. What is absolute zero? How is this point selected? 

11. State the kinetic theory and Avogadro's hypothesis. 

12. Solve problems 6, 7, 8, 9, and 10, pages 62 and 63. 

13. Answer questions 2, 4, 5, 9, and 10, page 51. 

14. Answer questions 1, 2, 3, 4, and 5, page 62. 

LESSON IV 
WATER; THREE STATES OF MATTER 

Assignment: Chapters VI and IX, McPherson and Henderson 

Occurrence. Water is certainly the most important compound 
we have to study. It occurs in greater quantity than any other. 
As the most widely used solvent it makes possible the occurrence 
of large numbers of chemical changes. It is necessary to the 
growth of plants and the digestion of food by animals. Its 



ELEMENTARY CHEMISTRY 19 

industrial uses are numerous. From a scientific point of view it is 
one of our most stable substances and a determination of its com- 
position is the starting point in developing the theories of chem- 
istry. Besides occurring on the earth's surface, it is found as a 
part of many substances. All plant and animal bodies contain it 
— nearly 70 per cent of the human body is water. Most foods 
contain a large amount of water and some mineral bodies are 
found to contain it, such as blue vitriol and gypsum. 

Impurities in water. Natural waters contain two kinds of 
impurities. They are mineral and organic. 

As the water travels over or through the earth it dissolves a 
portion of the solid matter with which it comes in contact. 
These mineral substances are usually common salt and compounds 
of calcium, magnesium, and iron. Since these substances prevent 
the formation of a lather with soap, such waters are called hard; 
water which lathers well contains little mineral matter and is 
called soft. These impurities do not make the water unfit to 
drink, unless present in very large amounts. They are harmful 
for industrial uses such as in laundries and boilers. In the latter 
case they form boiler scale. A more complete discussion of hard 
water and boiler scale will be given in Lesson XXL 

Organic impurities found in water consist of substances 
formed by the decay of vegetable and animal matter, substances 
found in sewage, and forms of living microorganisms which 
usually accompany such products. 

Effect of organic matter upon health. The effect of the 
organic matter present upon the health is very small. The harm 
is caused by the microorganisms, or bacteria, which are disease 
producing. Bacteria are forms of plant life. Not all bacteria 
produce disease; many kinds are not only harmless but useful, if 
not necessary, to human life; but some kinds produce serious 
diseases, such as tuberculosis, pneumonia, measles, smallpox, 
typhoid fever. 

As an example, take typhoid fever. This is a disease of the 
small intestine. It is here that the typhoid bacteria lodge and 
carry on their life processes, resulting in the fever. This disease 
can only be acquired by taking these bacteria into the alimentary 
canal by means of the mouth. This means that our food and 



20 ELEMENTARY CHEMISTRY 

drink must contain the germs of the disease. The typhoid bac- 
teria pass from the person suffering with the disease in the bowel 
excrement. Sewage containing this excrement may mix with the 
water supply, rendering it unfit to drink. Flies are also guilty of 
acting as carriers. They go from sewage deposits to food sup- 
plies and carry the bacteria on their feet. 

Detection of impurities. Mineral matter is detected by 
evaporation and analysis of the residue. Organic matter is detected 
by a sanitary analysis which shows the presence not only of organic 
matter itself but also of nitrogen compounds that result from the 
decay of organic matter. Sodium chloride always accompanies 
sewage, so its presence is significant. Large amounts of these 
substances make the water suspicious. 

A bacterial examination for typhoid bacteria is not practical. 
To be safe we would have to examine every drink of water. We 
must assure ourselves that the water is safe from contamination 
with sewage. A bacterial test to show this is made for colon 
bacteria, which are always present in the large intestine in great 
numbers and, therefore, always found in sewage. 

Purification of water. Distillation of water is the most 
effective way of purifying water. In this method the water boils 
and the steam condenses. The result is pure water. Solid matter 
is left behind. For most industrial purposes this is too costly, but 
it is used by the chemist and druggist. Artificial ice is made 
from distilled water. 

For drinking purposes boiling the water and allowing it to 
cool renders it safe. The heat kills the bacteria but does not 
remove the dissolved matter. On a large scale water is purified 
by cities by methods of filtration. Slow sand filters and so-called 
mechanical filters are used. 

Properties of water. When pure, water is colorless (except 
that it has a bluish tinge in thick layers), odorless, and tasteless. 
The temperatures of boiling, freezing, and greatest density are 
important. These are 100°, 0° and 4° C, respectively. A remark- 
able, and very useful property, is its ability to dissolve other 
substances. Such solutions are employed by chemists in study- 
ing the action of substances, since action is more rapid in 
solution. 



ELEMENTARY CHEMISTRY 21 

Water is a very stable substance; that is, it is hard to 
decompose. Heated to 2500° C, only a small amount is decom- 
posed into its elements. Its decomposition by means of the 
electric current is an indirect action. Water holds its oxygen 
firmly and is not a good oxidizing agent; however, carbon and 
some metals, as iron, will take the oxygen and set free hydrogen. 
In this way an illuminating gas is made from carbon and steam. 

y . Water forms hydrates with many compounds by combining 
directly, with them. The water in hydrates is called water of 
hydration. 4 Such water is in a state of chemical combination with 
the rest of the substance, as is shown by the fact that the per- 
centage composition is constant. For example, blue vitriol is a 
hydrate of copper sulfate and always contains exactly 36.03 per 
cent of water. Water in this state of chemical union can be 
easily removed from the compound by heating slightly above the 
boiling point of water. This shows that the chemical union is 
very weak compared with that between the elements in ordinary com- 
pounds, such as copper oxide, carbon dioxide, and even mercury oxide. 

Composition of water. The quantitative composition of 
water can be shown by methods of analysis or synthesis. Synthe- 
sis means building up from smaller parts. In analysis water is 
decomposed, as with the electric current, and the amounts of 
hydrogen and oxygen determined. In general, this shows 2 vol- 
umes of hydrogen to 1 of oxygen or, by weight, 1 part of hydrogen 
to nearly 8 parts of oxygen. The method is not very accurate. 

Synthetic methods are more accurate. Two of these have 
been used. That of Berzelius and Dumas was first used. It con- 
sists of passing hydrogen over hot copper oxide, with absorption 
tubes to collect and weigh the water formed. The loss in weight 
of the copper oxide is the weight of oxygen used. The weight of 
the water minus the weight of the oxygen gives the weight of the 
hydrogen. The method of Morley is more accurate. He used the 
eudiometer. This is filled with a mixture of hydrogen and oxygen 
which is exploded by the electric spark. The volumes of hydrogen 
and oxygen entering into the action can be calculated as weights. 
Morley's experiments gave a result of 1 part by tveight of hydrogen 
combining with 7.94 parts by weight of oxygen. This value is very 
important and should be remembered. 



22 ELEMENTARY CHEMISTRY 

If this experiment is carried out at ordinary temperatures, the 
product of the action is water in liquid form and its volume is 
insignificant. If, however, the eudiometer is surrounded by a 
steam jacket to keep the temperature at the boiling point, we 
find that 2 volumes of hydrogen combine with 1 volume of oxygen to 
form 2 volumes of steam. This value should be remembered, as 
use will be made of it later. The student should here read the 
experiments described on pages 74-79 inclusive. 

Hydrogen peroxide. Hydrogen and oxygen form another 
compound besides water. This compound contains 15.88 parts by 
weight of oxygen in combination with 1 part of hydrogen. Hence 
it is called hydrogen dioxide sometimes, which means twice as 
much oxygen as in water. It is not made by direct union of the 
elements. Barium peroxide and sulfuric acid will react to form 
barium sulfate and hydrogen peroxide. The barium sulfate is 
insoluble and may be filtered off, leaving a solution of hydrogen 
peroxide. 

Pure hydrogen peroxide is hard to obtain and is very unstable, 
breaking down with explosive violence into water and oxygen. 
Catalytic agents increase the speed of this decomposition. Because 
it gives off its extra oxygen easily, it is a good oxidizing agent. 
As such it is used as an antiseptic and bleaching agent; it oxidizes 
certain colored compounds and will bleach hair and some dyes. 

Law of multiple proportions. Considering the compounds 
hydrogen peroxide and water, we see that two elements may unite 
to form two compounds, but that they do so in different propor- 
tions. It is further observed that the quantity of oxygen which 
unites with a fixed weight of hydrogen is exactly twice as much in 
one case as in the other. Many other examples of the same kind 
could be named. These facts may be stated as the law of multi- 
ple proportions as follows: When two elements combine to form two 
or more compounds there is always a small whole-number ratio 
between the several weights of one element ivhich are in combination 
with a fixed weight of the other. This is the fourth law of chemical 
combination stated thus far. The fifth will be discussed in the 
next lesson. 

States of matter. We have learned that water exists as a 
solid, liquid, and gas. These states are known as the physical 



ELEMENTARY CHEMISTRY 23 

states of matter. Many other kinds of matter can exist in all 
three states. All gases have been liquefied and solidified, but all 
solids have not been liquefied or vaporized. A change from one 
state to another is a physical and not a chemical change. 

Liquids evaporate at all temperatures but the evaporation is 
more rapid as the temperature rises. The vapor thus formed 
exerts a pressure. This pressure is called the vapor pressure of the 
liquid. When this pressure is equal to the atmospheric pressure, 
the change to the gaseous state is rapid. A change in the atmos- 
pheric pressure would change the temperature at which the vapor 
pressure would overcome the atmospheric pressure. The boiling 
point of a liquid is defined as the temperature at which the vapor 
pressure just exceeds the atmospheric pressure. At this point a cer- 
tain amount of heat will be used up to change the liquid to a gas. 
This is known as the heat of vaporization. For water, 539 calories 
are required to change 1 gram of liquid to 1 gram of vapor at the 
boiling point. 

Ozone. A form of matter is known which can be shown to 
contain only oxygen but which has properties quite different from 
those of oxygen. The main difference is that ozone is much more 
active and has a much higher energy content. It changes to 
oxygen; consequently it acts as a very powerful oxidizing agent. 
It is a bleaching agent, and it destroys many low forms of life 
and is therefore used as a disinfectant for purifying water and air. 

It can be prepared by passing a silent electric discharge 
through oxygen. Energy is thus stored up in the ozone as chemi- 
cal energy. Several forms of an element differing only in the 
chemical energy they contain are said to be allotropic forms of 
that element. Graphite and diamond are allotropic forms of 
carbon. 

EXERCISES 

1. Discuss the occurrence and uses of water. 

2. What are the kinds of impurities in water? State the harm each does. 

3. How are these impurities detected? 

4. Discuss the methods for purifying water. 

5. State the properties of water. 

6. Give the weight composition and the volume composition of water. 
What volume of steam would be formed from 2 volumes of hydrogen? i 

7. Describe methods by which this composition can be determined. 



24 ELEMENTARY CHEMISTRY 

8. State composition, preparation, properties, and uses of hydrogen 
peroxide. 

9. State the law of multiple proportions and give an example of it. 

10. What is meant by vapor pressure of a liquid, boiling point, heat of 
vaporization, allotropic? 

11. State the composition, preparation, properties, and uses of ozone. 

12. Solve problems 14, 15, 16, and 17, page 115. 

13. Solve problems 12, 13, 14, 15, and 16, page 85. 

14. Answer questions 1, 2, 4, 5, 6, and 7, page 84. 

15. Answer questions 2, 3, 4, 6, 7, 10, 11, 12, and 13, page 115. 

LESSON V 

ATOMIC THEORY; EQUATIONS 
ATOMIC THEORY 

Assignment: Chapter VII, McPherson and Henderson 

Law of combining weights. Of the laws previously studied 
three deal with the weight relations in chemical change. These 
are (1) the law of conservation of matter, (2) the law of definite 
composition, and (3) the law of multiple proportion. The law of 
combining weights is the fourth. These four laws are the funda- 
mental laws of weight relations in chemical change. 

In the study of water it was found that 1 gm. of hydrogen 
combined with 7.94 gms. of oxygen. This compound and its com- 
position furnish the most satisfactory numerical relation to use for 
a starting point in studying other numerical relations. If we study 
calcium oxide, we find that 7.94 gms. of oxygen combine with 
19.88 gms. of calcium; 19.88 gms. of calcium combine with 16 
gms. of sulfur. We may, by studying the composition of a large 
number of compounds of different elements, that is, by analyzing 
their compounds, make a list of the elements and assign a number 
to each element. Such number will represent the weight of the 
element to which it is assigned that will combine with the 
assigned weight values of the other elements. Using whole num- 
bers for the values, a brief list is here given: H = l, 0=8, Ca = 20, 
S = 16, K = 39, CI = 35, Br = 80, Na=23, 1 = 127, Cu=63, F = 19, 
Zn = 65. 

The student must keep in mind that these numbers are 
experimentally determined and are facts beyond question. Not 
every element in the above list will combine with every other, but 



ELEMENTARY CHEMISTRY 25 

where they do not combine they will take the place of each other 
in compounds in weight proportions corresponding to the above 
numbers. Thus, hydrogen does not combine with zinc, but in 
their compounds with chlorine 1 gm. of hydrogen will be equiva- 
lent to 65 gms. of zinc. These weights are called combining, or 
equivalent, weights. 

The standard. These numbers are only relative and, for the 
above list, are the weights found to combine with or be equivalent 
to 1 gm. of hydrogen. They are relative to hydrogen. Any other 
element could be taken as a standard for comparison and any 
number assigned to it. The ratio between the combining weights 
would not change though the numbers would be different if some 
other number was assigned to hydrogen. 

As a matter of fact, oxygen is most commonly used as the 
standard, because more elements will combine with it than with 
hydrogen. If 1 gm. of hydrogen combines with 7.9-1 gms. of 
oxygen, then 1.008 gms. of hydrogen will combine with 8 gms. of 
oxygen. This last value is the one selected for a standard. The 
combining, or equivalent, weight of an element is then the weight of it, 
in grams, which will combine with 8 gms. of oxygen. Since some 
elements combine with oxygen in more than one proportion they 
have more than one combining weight. The numbers which rep- 
resent the several combining weights of any element are always 
small whole-number multiples of the smallest combining weight 
(law of multiple proportion). 

The law of combining weights is stated as follows: To each 
element may be assigned a number which in itself, or when multiplied 
by some small integer, expresses the weight by which the element com- 
bines with other elements. 

Explanation demanded. These four laws state the facts of 
chemical imion, but to the normal mind something is still lacking. 
Why this regular behavior? This question demands an answer. 
The law of multiple proportions, in particular, presses for an 
explanation. The answer must lie in the nature of matter. The 
answer given to this question by Dalton is the simplest and most 
satisfactory. 

There are five oxides of nitrogen. With 7 gms. of nitrogen 
we find combined respectively 4, 8, 12, 16, and 20 gms. of oxygen. 



26 ELEMENTARY CHEMISTRY 

Why always a whole-number multiple of 4, as 2, 3, 4, 5 times 4? 
Never do we have a fractional multiple, as 3.5 or 4.2. Does it not 
suggest to us, that the amount of oxygen represented by 4 parts 
by weight represents the unit of oxygen by which chemical action 
is taking place, and that in the other compounds we have 2, 3, 4, 
or 5 of these units? This is what Dalton thought when he stated 
the atomic theory. 

Atomic theory. This theory suggests that the laws of chemi- 
cal union are explained, if we assume that all elements are made 
up of small unit particles, which we call atoms. The atom is 
defined as the smallest particle of an element that can take part in 
chemical changes of that element. All the atoms of each element 
have the same weight, but atoms of different elements have differ- 
ent weights. When two or more elements unite to form a chemi- 
cal compound, it requires a definite number of atoms of each ele- 
ment; hence, a compound always has the same composition. The 
weight of the atoms do not change in chemical change because 
there is no gain nor loss in total weight. 

If two elements form more than one compound, it is because 
one, two, three, or four atoms of one element combine with a 
fixed number of atoms of the other element; thus establishing the 
small whole-number ratio between the weights of one element 
which combine with a fixed weight of the other (multiple proportion) . 

If these atoms always have the same weight, then the num- 
bers assigned as the combining weights of the elements, or some 
small multiple of them, can be taken to indicate the relative 
weights of the atoms themselves. These, then, are the atomic 
weights. The actual weights of the atoms cannot be determined. 
We must use their relative weights as referred to some one of the 
elements as a standard. Oxygen is the most satisfactory to use as 
a standard. The ratio of hydrogen to oxygen in water is the 
fundamental ratio. 

It will be shown in Lesson XVI that water contains two 
atoms of hydrogen and one of oxygen. We must either use 0.5 as 
the atomic weight of hydrogen or double the numbers in the ratio 
1.008 to 8 for the composition of water. If we do the latter, we 
have 2.016 to 16. The atomic weight of hydrogen can then be taken 
as 1 . 008 and that of oxygen as 16. This is the standard generally used. 



ELEMENTARY CHEMISTRY 27 

Molecules. The atom is tJw smallest unit of matter that enters 
into chemical changes; the smallest unit that takes part in physical 
changes is called the molecule. There are molecules of elements as 
well as of compomids. Usually a molecule is composed of two or 
more atoms. The molecules of compoimds are made up of the 
atoms of the elements composing them. The molecules of ele- 
ments are made up of the atoms of the single element, frequently 
two or more, but some elements are known which contain but 
one atom in a molecule, for example, mercury. 

EQUATIONS 
Assignment: Chapter VIII, McPherson and Henderson 

Formulas and percentage composition. The analysis of a 
compound is usually expressed in percentage, or parts per hun- 
dred. The chemist can express composition in terms of atomic 
weight quantities in the molecule, or the number of atoms of each 
element in the molecule of the compound. For this purpose the 
symbol of the element is used to represent one atomic weight of 
the element. Thus Hg stands for 200 parts by weight of mercury, 
and O for 16 parts by weight of oxygen. The compound mercuric 
oxide is represented by HgO. This is called a formula. 

If we know the formula of a compound and the atomic 
weights of its elements, it is a simple problem to calculate the 
parts per hundred. If the atomic weight of mercury is 200 and 
that of oxygen is 16 and one atom of each combine to form a 
molecule of the compound, this molecule can be said to have a 
weight of 216 relative to our standard. Divide 200 by 216 and 
we have the per cent of mercury in the compound. 

Determination of formulas. On the other hand, if we know 
the percentage composition and the atomic weights, we can cal- 
culate the number of atoms in a molecule of the compound, that 
is, determine its formula. To determine the actual molecular 
formula it is necessary that the molecular weight of the compound 
be known. It is not always possible to know this. We may still 
determine the formula that represents the smallest possible per- 
centage composition of the compound. This is called the empiri- 
cal, or experimental, formula. 



28 ELEMENTARY CHEMISTRY 

Sulfuric acid contains 2.06 per cent H, 32.69 per cent S, and 
65.25 per cent 0, that is, that many grams in 100 gms. If we 
divide these numbers by the atomic weights of their elements, we 
will obtain the relative number of atoms of each element in the 
compound. These relative numbers are quite likely to be frac- 
tions, but atoms do not exist as fractions; they are present only 
in whole quantities. Keep in mind that the relative numbers of 
atoms are a ratio and their value can be changed by dividing by a 
common factor without changing the ratio. The simplest formula 
will be one that contains but one atom of at least one element. 
The relative numbers of the different kinds of atoms in the for- 
mula would be found by dividing the fractional values first 
obtained by the smallest of the lot. Thus: 

2.06-4-1.008 = 2.04] 
32.694-32.06 = 1.02 1 4- 1.02 
65.25-M6 =4.08j 

Therefore sulfuric acid contains two atoms of hydrogen, one of 
sulfur and four of oxygen. This is written H 2 S0 4 . In a formula 
the number of atoms is represented by a small figure written 
slightly below the line and after the symbol to which it refers. 
The actual molecular formula may be some multiple of this, as 
H 4 S 2 8 or H 6 S 3 12 . 

Facts expressed by formulas. Just as the symbol represents, 
not simply an abbreviation of the name of an element, but its 
atomic weight, so the formula of an element or compound shows 
more than simply the kind of elements in its composition. In 
addition to showing the number and kind of atoms in a molecule 
of a compound, the formula shows the relative weights of these 
atoms and the molecular, or formula, weight of the compound. 
Thus, H 2 S0 4 shows sulfuric acid to contain two atomic weights of 
hydrogen, one atomic weight of sulfur, and four atomic weights of 
oxygen. The weight of the molecule of sulfuric acid is, therefore, 

(2Xl.008)+32.06+(4xl6) =98.076 

This is called the molecular weight because this weight represents 
the simplest composition of sulfuric acid. 

In dealing with the substance we use grams or pounds. We 
do not know the gram weight of atoms. Since the atomic 



ELEMENTARY CHEMISTRY 29 

weights are relative to each other, the ratio is the same, if we 
allow H to represent 1.008 gms. of hydrogen; S, 32.06 gms. of 
sulfur; O, 16 gms. of oxygen; and H 2 S0 4 , 98.076 gms. of sulfuric 
acid. These weights are called the gram-atomic weights for the 
elements and the gram-molecular, or formula, weights for the 
compounds. 

Equations. A grouping together of the formulas of the sub- 
stances entering into a reaction and those of the products formed 
and connecting the two groups by a convenient sign constitutes a 
chemical equation. Such an equation is not mathematical in its 
nature. It is simply a shorthand method of representing the 
facts of the chemical change which has taken place. 

If experiment shows that 98 gms. of sulfuric acid will act 
with 65 gms. of zinc to give 161 gms. of zinc sulfate and 2 gms. of 
hydrogen, we can briefly express these facts by using the formulas 
that represent the weights of the compounds. Thus, 

H 2 S0 4 -f-Zn-*ZnS0 4 +H 2 

Such an equation would be read as follows: One formula weight of 
sulfuric acid and one atomic weight of zinc react to give one for- 
mula weight of zinc sulfate and one formula weight of hydrogen. 
There are two atoms in the molecule of hydrogen. 

In order to write an equation it is necessary to know certain 
facts about the chemistry of the reaction. Only then can you 
write the equation. In the above example it must be shown by 
experiment that 98 gms. of H 2 S0 4 react with 65 gms. of zinc to 
give 161 gms. of ZnS0 4 and 2 gms. of hydrogen. This experi- 
mental work is not possible for the beginning student. He 
must, however, get the same facts from the textbooks or an 
instructor. 

Classes of equations. A study of chemical changes shows 
that practically all may be represented by four types of reactions: 

1. Combination, in which two elements or compounds unite to form one 
compound: 

2H 2 +0 2 — >2H 2 
2Hg+0 2 ->2HgO 
Fe +S ->FeS 

c +o 2 ->co 2 

CaO+H 2 0->Ca(OH) 2 



30 ELEMENTARY CHEMISTRY 

2. Decomposition, in which a compound decomposes into simpler parts: 

2KC10 3 ->2KC1 +30 2 
2H 2 -->2H 2 +0 2 
CaC0 3 ->CaO +C0 2 
2HgO -»2Hg -f0 2 

3. Substitution, in which one element replaces another in a compound: 

2Na+2H 2 ->2NaOH+H 2 
Zn +H 2 S0 4 ->ZnS0 4 +H 2 
Fe +2HCl->FeCl 2 +H 2 

4. Double decomposition, in which two compounds react, the molecules 
of both decomposing and recombining to form two new compounds: 

Ba0 2 +H2SO4 — >BaS0 4 +H 2 2 
AgN0 3 +NaCl ->AgCl +NaN0 3 
HC1 +NaOH->HOH -fNaCl 
H 2 S0 4 +2NaCl->Na 2 S0 4 +2HCl 

Heat and conditions of a reaction. The above equations 
represent the products of the reaction and the proper quantities of 
each substance to complete the reaction, but they do not represent 
the energy change that always takes place. In many cases this 
energy change is a heat change. Heat is either absorbed or given 
off by the action. While the equations do not show this, the 
student must not forget that such a change is always occurring. 
Neither do the equations show the conditions under which the 
action will occur. These conditions are important, and the 
proper ones must be obtained to get the reaction. 

The equation is a statement of the relative weights of the 
reacting substances in terms of their formula weights. If the 
equation correctly represents the facts of the reaction, the same 
ratio between the weights of the several substances holds good, no 
matter in what weight units they may be measured — grams, 
pounds, ounces, or what not. If we know the gram or pound 
weight of one substance involved in the reaction, we can use the 
ratio of relative weights shown by the formula weights in the 
equation to calculate the actual weights of the other substances 
concerned. The equation must always be written to show this 
relation before a start can be made to solve a problem. 

Examples. How many grams of oxygen will be required to burn 50 
grams of carbon? 

The equation for this action is C+0 2 — >C0 2 . This shows that 12 parts 
by weight of carbon combine with 32 parts by weight of oxygen to form 44 



ELEMENTARY CHEMISTRY 31 

parts of carbon dioxide. If 12 gms. of carbon react with 32 gms. of oxygen, 
1 gm. of carbon will react with 32-V-12, or 2.66 gms. of oxygen, and 50 gms. 
with 50X2.66 gms., or 133.3 gms. of oxygen. Or, the problem may be ex- 
pressed in the form of a proportion: 

12: 32:: 50: X 

32X50 

X = = 133.3 gms. 

12 & 

In solving a problem, the formula weights of the two substances involved 
in the solution should be set down under the formulas of the compounds. 
Thus, if it is required to find how many grams of H2O2 can be obtained from 
40 grams of Ba0 2 

Ba0 2 +H2SO4— >BaS0 4 +H 2 02 
169.37 34.016 

Then 

169.37: 34.016:: 40: X 
Or 

169.37 _ 40 

34.016" X 

„ 34.016X40 



169.37 



= 8.03 gms. H 2 2 



EXERCISES 

1. Define atom, molecule, symbol, formula, equation, formula weight, 
gram-molecular weight. 

2. State the four fundamental laws of chemical action treating of mate- 
rial changes. 

3. Give a full statement of the atomic theory. 

4. Give names of four types of chemical reactions. Define each. Give 
two equations to illustrate each. 

5. Solve problems 2, 3, 4, 5, 7, 8, 9, 13, and 14, pages 102-103. 

SEND EXERCISES FOR LESSONS I=V TO THE SCHOOL 



LESSON VI 
CARBON AND CARBON DIOXIDE 

Assignment: Chapter X, McPherson and Henderson 

Occurrence of carbon. Carbon is a familiar element, occur- 
ring in impure form in coal and charcoal. Its compounds are 
widely distributed and are commonly met with in our daily life. 
Carbon, together with hydrogen, is a necessary constituent of all 
of that large class of compounds known as organic. By organic 
compounds was originally meant those associated with living mat- 



32 ELEMENTARY CHEMISTRY 

ter, plant or animal. That meaning is no longer satisfactory as 
many organic compounds have been made without association 
with living things. 

Yet the composition of all living matter very largely consists 
of these so-called organic compounds, or compounds of carbon and 
hydrogen and their derivatives. Such compounds are also found 
in natural gas, petroleum, and the distillation products of wood 
and soft coal. Limestone and marble are carbonates and are salts 
of an acid of carbon, just as the chlorides are salts of an acid of 
chlorine. 

Forms of carbon. The element carbon is capable of existing 
in three allotropic forms. Two of these are crystalline and one 
amorphous. The crystalline forms are diamond and graphite. 
The amorphous form is pure charcoal made from sugar. A com- 
plete statement of the properties and preparation of diamond and 
graphite will be found on pages 117 and 118. 

That these substances are allotropic, that is to say, that they 
consist of the same kind of matter but differ as to the amount of 
energy contained, is shown by experiments in the combustion of 
them. If 1 gm. of each be burned, the same weight of carbon 
dioxide will be obtained but different amounts of heat energy will 
be set free from the several forms. 

Natural varieties of amorphous carbon. Several varieties of 
amorphous carbon are well known; the differences in properties are 
due to impurities or to variation in physical structure resulting 
from the methods of formation or manufacture. Some occur in 
nature as the various forms of coal. Anthracite, or hard, coal is 
chiefly elementary carbon. Soft, or bituminous, coal contains a 
considerable part of the carbon combined with other elements, 
chiefly hydrogen. Oxygen, nitrogen, and sulfur are found also. 

Coal deposits are considered to be the result of a process of 
destructive distillation occurring in nature. In a very early geo- 
logical age of the earth's history the atmosphere contained a much 
larger percentage of carbon dioxide than it does at present. Car- 
bon dioxide is a necessary plant food. In that age there is evi- 
dence that a very rank and luxurious form of vegetation existed. 
In the course of successive changes in the earth's surface this 
heavy vegetation became covered over with earthy deposits. 



ELEMENTARY CHEMISTRY 33 

Thus confined, the heat of the earth carried on a decomposition 
which resulted in the formation of deposits of coal containing 
varying amounts of elementary carbon. 

The older formations are anthracite, the later formations 
bituminous coal. Natural gas is a product of this process of 
natural distillation. Lignite, a substance closely related to 
bituminous coal, is a still younger formation and contains only a 
small amount of carbon as the element. Peat is another form of 
vegetable matter in the early stages of decomposition. It is 
largely mixed with earthy matter but, when dried, makes a satis- 
factory fuel. 

Artificial varieties of amorphous carbon. Other forms of 
amorphous carbon can be manufactured from carbon compounds. 
Coke is made by the destructive distillation of bituminous coal, 
that is, by the heating of the coal in a retort or furnace from 
which air is excluded, so that combustion does not take place. 
The coke consists of carbon and the mineral matter of the coal. 
Other products of the heating of bituminous coal in the absence of 
air are coal gas, ammonia, and coal tar. All are valuable and will 
be considered fully in later lessons. 

Wood charcoal is prepared from w T ood in a way similar to that 
for making coke from coal. Among the other products of the 
destructive distillation of wood are wood alcohol and acetic acid. 
Lampblack is a product of the destructive distillation of oils and 
gas containing high percentages of carbon. Bone black, or animal 
charcoal, is the product of the destructive distillation of bones and 
refuse animal matter. 

Varieties cf amorphous carbon not allotropic. These several 
forms of carbon are simply varieties of amorphous carbon and not 
allotropic modifications of the element. An experiment in com- 
bustion w 7 ill show this to be the case. 

Since these forms are impure, such amounts must be taken for 
combustion as will contain the same amount of carbon. This can 
be done by taking such amounts of the substances as will give on 
combustion equal weights of carbon dioxide. If the different forms 
were allotropic, each would generate on combustion a different 
amount of heat energy. In this experiment, how r ever, it would 
be found that for equal weights of carbon dioxide formed equal 



34 ELEMENTARY CHEMISTRY 

amounts of heat energy would be liberated. Therefore, they are 
not allotropie, since their carbon contains the same energy. 

Properties and uses. Many of the physical properties of car- 
bon vary with the different allotropie forms. All, however, are 
odorless, tasteless, insoluble in water, and volatile only in the 
intense heat of the electric arc. Bone black, especially, has the 
property of absorbing colors and odors, and liquids are purified 
and decolorized by being passed through bone-black filters. 

Carbon is used in gas masks because of this property of 
absorbing gases. During the War it was so used. A special form 
of carbon for this purpose was prepared from coconut shells, 
peach and prune seeds, etc. Cane-sugar solutions are decolorized 
prior to the crystallization of the granulated sugar by being 
filtered through carbon. 

In its chemical behavior carbon is characterized by its inac- 
tivity. It does not act on the ordinary acids except those that are 
strong oxidizing agents. At high temperatures it will combine with 
oxygen, sulfur, and hydrogen. With some metals it forms carbides. 

Its chief use is as a fuel. However, large amounts are used as 
a reducing agent because of its great tendency to combine with 
oxygen. In obtaining metals from their oxides carbon serves as 
the reducing agent; in the blast furnace carbon reduces the iron 
oxide to iron; other examples are the reduction of zinc and tin 
oxides with carbon. 

Carbon dioxide. There are two well-known oxides of carbon: 
carbon dioxide (C0 2 ) and carbon monoxide (CO). Only the for- 
mer will be studied at this time. Carbon dioxide occurs in the 
atmosphere, about 3 parts in 10,000. It is a colorless gas formed 
by the combustion in oxygen of substances containing carbon, 
such as the ordinary fuels; it is also a product of respiration, 
being formed by the oxidation of foods in the body; and fermenta- 
tion and decay of organic matter likewise form it. Its presence in 
the atmosphere is essential to plant life. 

For many purposes it is prepared by the action of an acid on 
a carbonate. Calcium carbonate (CaC0 3 ) is usually used because 
of its common occurrence. The following equation represents the 
action when hydrochloric acid is used: 

CaC03+2HCl->CaCl 2 +H 2 0+C0 2 



ELEMENTARY CHEMISTRY 35 

For a description of the method see page 123. It may also be 
prepared by heating limestone as in the preparation of lime 
(CaO). It has been prepared commercially by fermentation in 
connection with the manufacture of malt beverages. 

Properties and uses. Carbon dioxide is a very heavy gas 
and can be poured like water. For this reason it sometimes col- 
lects in wells and mines. It is 1.5 times as heavy as air. It is 
colorless and odorless. It can be liquefied and solidified; the 
solid is obtained by the cooling effect produced by the sudden 
expansion of the substance when released from pressure (page 125). 
Solid carbon dioxide has a low enough temperature to freeze 
mercury; it is used in surgery for freezing local areas in minor 
operations. 

Carbon dioxide cannot be burned, as it contains all the 
oxygen that carbon is capable of holding in combination. Neither 
will it support combustion as it is a very stable compound. The 
word "stable" is used in the sense of not being easily separated 
into its elements. It holds its oxygen in firm combination and 
cannot give up any to support the combustion of other sub- 
stances. 

Its presence in the atmosphere can be shown by its action 
with calcium hydroxide, or limewater, (Ca(OH) 2 ), a white precipi- 
tate being formed when the gas is passed into limewater. The 
chemical change is represented as follows: 

Ca(OH) 2 +C0 2 ->CaC0 3 +H 2 
Carbon dioxide is a plant food and is absorbed from the air by 
the leaves. 

It is the gas that causes effervescence in many beverages, 
such as soda water, mineral waters, champagne, and beer; the gas 
is held in solution under pressure and escapes when the cork is 
removed. Its use as a fire extinguisher is important. In all port- 
able fire extinguishers, or chemical fire engines, as they are some- 
times called, a bottle of concentrated sulfuric acid is upset or 
broken into a solution of sodium carbonate. The gas formed 
creates pressure to force the saturated solution of carbon dioxide 
out of the generator (see page 126). 

Carbon dioxide is not a poison, though persons could not live 
in an atmosphere of it since it would exclude oxygen from the 



36 ELEMENTARY CHEMISTRY 

lungs and death would result from lack of oxygen, as it does in 
cases of strangulation. Other compounds of carbon will be 
studied in Lessons XVII and XVIII. 

EXERCISES* 

1. State the occurrence of carbon in nature. 

2. Name the allotropic forms and varieties of carbon. 

3. Give an experiment which will show that diamond and graphite are 
allotropic. 

4. Give one to show that anthracite coal and coke are not allotropic. 

5. How may diamonds be made artificially? 

6. Give the sources and properties of diamonds. 

7. Give the sources and preparation of graphite. 

8. Give the properties and uses of graphite. 

9. How is pure charcoal made? 

10. Discuss the sources, origin, and properties of various kinds of coal. 

11. Give methods for making coke, charcoal, bone black, and lamp- 
black. 

12. State the physical properties and chemical conduct of carbon. 

13. Give the uses of carbon. 

14. Where is carbon dioxide found, and how is it formed by nature? 

15. Give the methods for preparing carbon dioxide and the equations 
where possible. 

16. How is solid carbon dioxide made? What are its uses? 

17. Give the properties and the chemical conduct of carbon dioxide. 

18. What are the uses of carbon dioxide? How is it used as a fire 
extinguisher? 

19. Solve problems 7, 12, 15, and 16, page 127. 

20. Answer questions 4, 5, 9, 11, 13, and 14, page 127. 

LESSON VII 

NITROGEN; THE ATMOSPHERE 
NITROGEN 
Assignment: Chapter XI, McPherson and Henderson 

Occurrence of nitrogen. This element is very plentiful, but 
not so widely distributed as some others. It occurs chiefly in the 
atmosphere, of which it constitutes 78 per cent by volume, this 
source seeming to be inexhaustible. 

Air was at one time thought to be a simple substance. 
Rutherford obtained nitrogen from air by burning charcoal and 

*Prepare this set of Exercises and hold it until those for Lessons VII, VIII, IX and X are 
also prepared and then send all five sets to the School. 



ELEMENTARY CHEMISTRY 37 

phosphorus in it. Lavoisier called the residue (nitrogen) azote, 
meaning unable to support life; the name "nitrogen" is used 
because this element is a characteristic constituent of niter, or 
saltpeter. The chemical name for saltpeter is potassium nitrate 
(KN0 3 ). Nitrogen also occurs in sodium nitrate (NaN0 3 ), found 
in Chili and called Chili saltpeter. Many organic compounds con- 
tain nitrogen, for example, those that compose the substances 
known as proteins. Proteins are an important class of foods as 
they are the tissue builders. The human body contains 3 per 
cent nitrogen. 

Preparation. For commercial purposes nitrogen is prepared, 
along with oxygen, by the liquefaction of air. The nitrogen, 
having a lower boiling point, boils off first and can be collected. 
Liquid air will be considered later in this lesson. Nitrogen pre- 
pared in this way contains some oxygen but serves for those 
industrial purposes in which nitrogen is one of the raw materials. 
Nitrogen is obtained from the air by other methods based upon 
the action of some substance which will combine with the oxygen 
and leave the nitrogen. Phosphorus and copper are used (see 
page 129). 

None of these methods gives pure nitrogen. There are 
several rare elements in the air which are left along with the nitro- 
gen when the oxygen is removed. As they are more inert chem- 
cally than nitrogen, they do not interfere with its use for many 
purposes. Pure nitrogen can be prepared only from compounds 
which contain it. Ammonium nitrite is most often used. This 
compound is so unstable that it cannot be obtained in the market, 
but has to be prepared fresh at the time of its use; its decompo- 
sition as fast as it is formed produces the nitrogen. The two 
steps are shown by the following equations: 

NaN0 2 +NH 4 Cl^NaCl+NH 4 N0 2 
NH 4 N0 2 ->2H 2 0+N 2 

Properties. In its chemical behavior nitrogen is less active 
than all the elements previously studied and much less active than 
oxygen and hydrogen. It shows no tendency to combine with 
other elements at ordinary temperatures. It combines with mag- 
nesium, oxygen, hydrogen, and a few other elements with diffi- 



38 ELEMENTARY CHEMISTRY 

culty and only at high temperatures. Conditions can be established 
under which nitrogen and hydrogen will combine to form ammonia 
on a commercial scale. The same is true for the reaction between 
oxygen and nitrogen. When ammonia and nitric acid are studied 
in Lesson XII, these reactions will be taken up again. 

The compounds of nitrogen with metals are called nitrides. 
Its compound with hydrogen is ammonia (NH 3 ), a common and 
useful substance. With oxygen it forms a number of oxides. 
These compounds will be studied later. Like carbon dioxide it is 
not a poison, but people could not live in an atmosphere of pure 
nitrogen because of the lack of oxygen. Its physical properties 
are stated on page 131. 

Uses. Nitrogen is used to fill thermometers and light bulbs 
and to prepare ammonia, nitric acid, and calcium cyanamide, 
which substances are used in making fertilizers (see page 462). 
Nitrogen is necessary to plant life as it is the element that builds 
up the protein tissues; but plants have no mechanism for absorbing 
free nitrogen, they can only absorb it when it exists in the form of 
water-soluble compounds which can be taken up from the soil by 
their roots. Nitrogen, to be available, must be converted into 
ammonium salts, nitrates, or compounds which on action with 
water will form one of these. Many efforts have been made to 
devise methods for fixing atmospheric nitrogen so that the supply 
of nitrogen fertilizer material might be increased. 

In this connection the action of such plants as clover, alfalfa, 
beans and peas, belonging to the family of legumes, is interesting. 
Indirectly they assimilate nitrogen. They have tubercles, or 
nodules, on their roots in which a form of bacterial life develops. 
These microorganisms are able to bring about the union of nitrogen 
and oxygen, forming compounds which form water-soluble nitrates 
with substances in the soil. Some of these nitrates are absorbed 
by the plant itself; others remain in the soil to increase its fertility 
(page 132). 

Rare elements. Argon, helium, neon, krypton, and xenon 
are rare elements occurring in the air in small amounts. They 
are similar in properties and noted for complete chemical non- 
activity. They form no known compounds. Argon is the most 
plentiful and helium the most important. Next to hydrogen, 



ELEMENTARY CHEMISTRY 39 

helium is the lightest element. During the late war considerable 
quantities of helium were discovered in the gas from certain oil 
wells in Texas. The quantity was great enough to make it pos- 
sible to use it to fill war balloons. It has the advantage over 
hydrogen of being non-combustible, which means that when the 
gas bag is punctured by an enemy shot, there is no combustion 
of the balloon but a slow leakage which enables the observer to 
come to the earth in safety. 

THE ATMOSPHERE 
Assignment: Chapter XII, McPherson and Henderson 

Composition. Air was formerly considered an element, but 
the obtaining of oxygen and nitrogen from it contradicts this idea. 
Besides these elements it contains water vapor and carbon dioxide 
in small amounts as well as the rare elements. Consult page 135 
for the percentage of each constituent. Water vapor varies but 
the amounts of the others are essentially constant (page 136). 

Of these constituents oxygen is necessary to support the life 
of animals, carbon dioxide to support the life of plants, and w^ater 
vapor to support the life of both, as it prevents evaporation of 
water from the tissues. Nitrogen is essential to dilute the others, 
especially oxygen, and so to diminish the intensity of its action. 

Quantitative analysis of air. The quantitative analysis of air 
consists in determining the amounts of each of the four principal 
constituents, using the reactions described on pages 136-137. In 
actual practice it is necessary to start with a measured volume 
of air and apply these reactions in a certain sequence. For exam- 
ple, air is passed through a gas meter to measure its volume, and 
then through a weighed vessel containing calcium chloride (CaCk), 
which is a good drying agent. The increase in the weight of the 
vessel is the weight of the water in the measured volume of air. 
The remaining air is then passed through a weighed vessel con- 
taining sodium hydroxide or calcium hydroxide, which absorbs 
the carbon dioxide: 

2NaOH+C0 2 ->Na 2 C0 3 +H 2 

The air is next passed over, or allowed to stand for a time in con- 
tact with, yellow phosphorus, which combines with oxygen. The 



40 ELEMENTARY CHEMISTRY. 

loss in volume measures the amount of oxygen. Finally the gas 
left may be passed over hot magnesium, when the nitrogen com- 
bines with it to form magnesium nitride (Mg 3 N 2 ), leaving only 
the rare elements of the air. 

Changes in composition. The average composition of the air 
remains nearly constant, but there are continual changes tending 
to increase or decrease oxygen, carbon dioxide and water vapor. 
Carbon dioxide is added by the exhaled breath of animals, the 
burning of carbon compounds (fuels), the decay of organic matter, 
and escapes from volcanoes. Carbon dioxide is removed by the 
action of plants, which absorb it through their leaves. Sunlight 
enables plants to utilize the carbon to build up the carbon com- 
pounds of their tissues. The weathering of rocks also absorbs 
carbon dioxide. 

Oxygen is added to the air by the action of plants. It is 
removed by the inhalation of animals in breathing and by the 
combustion of fuels. These changes, which offset each other, 
together with the mixing influence of the winds, keeps the composi- 
tion of air constant. 

Air a mixture. Two reasons for considering air to be a mix- 
ture are given on page 139: (1) The composition of air dissolved 
in water is not the same as that of undissolved air. (2) Air does 
not have a constant boiling point as a compound always has. A 
third reason may be given. The formation of a compound 
always is evidenced by some energy change. Nitrogen and oxy- 
gen when brought together, in the proportions for forming air, 
give a product which is not different from air, but there is no 
evidence of an energy change, such as evolution or absorption of heat. 

Impure air. In rooms and halls where people live or meet, a 
frequent renewal of the air is necessary to health. The bad 
effects of stale air arise from substances exhaled from the lungs. 
It is not supposed that the ill effects are due to an excess of car- 
bon dioxide or to a lack of oxygen in ordinary cases; they are due, 
rather, to an excess of moisture, high temperature, and the pres- 
ence of certain poisonous organic substances exhaled with the 
carbon dioxide. The air should be in constant slow motion, 
otherwise that near a person's body becomes saturated with 
moisture (page 140). 



ELEMENTARY CHEMISTRY 41 

Liquid air. The properties of liquid air are stated on page 
141. It is used in preparing oxygen and nitrogen for industrial 
purposes. The method of liquefaction is discussed on page 108. 
Dewar flasks are used for holding liquid air. They have a double 
wall with a vacuum between. The common thermos bottle is 
constructed on the same plan. 

Artificial ice. The principle used in making liquid air is that 
of compression and sudden expansion. The suddenly expanding 
gas absorbs great quantities of heat and is used to cool other por- 
tions of the substance below the liquefaction point. The same 
principle is applied in making artificial ice. Ammonia is com- 
pressed and allowed to expand through a small valve into pipes 
running into brine, or salt solution, which will not freeze until 
quite a low temperature is reached; as a result of the rapid 
expansion of the ammonia the brine becomes so cold that distilled 
water in tanks placed in it is cooled to the freezing point (page 
111). 

EXERCISES 

1. Give the sources of nitrogen. 

2. How is nitrogen obtained for industrial purposes? 

3. Describe the method for liquefying air. 

4. How are phosphorus and copper used to obtain nitrogen? 

5. Write equations and explain the method for making pure nitrogen. 

6. Give the physical properties of nitrogen. 

7. Discuss the chemical conduct of nitrogen. 

8. What are the uses of nitrogen? 

9. Explain the action of legumes on nitrogen. 

10. Name the rare elements. What is the use for helium? 

11. Give the composition of the air. 

12. How could this composition be determined quantitatively? 

13. Discuss changes in the composition of the air. 

14. Why is air believed to be a mixture? 

15. State the properties of liquid air. 

16. What is a Dewar flask; a thermos bottle? 

17. Describe the manufacture of artificial ice. 

18. Solve problems 9, 10, 11, 12, and 13, page 134. 

19. Solve problems 7, 8, 9, 10, and 11, page 142. 

20. Answer questions 1, 3, 4, and 5, page 134. 

21. Answer questions 2, 3, 4, 5, and 6, page 142. 



42 ELEMENTARY CHEMISTRY 

LESSON VIII 
SOLUTIONS; IONIZATION 

Assignment: Chapter XIII, McPherson and Henderson 

Definition. Most of us are familiar with the term solution, 
and there is a widespread tendency to think of it as referring to 
something liquid. Most solutions are of this kind, but some very 
common ones are solids, as glass and brass. However, take a 
common example, water and sugar; one substance is a liquid, the 
other a solid; on adding the sugar to the water, the solid crystals 
disappear and the particles of sugar can no longer be distinguished, 
even by the aid of a very high-power microscope. 

On the other hand, that change fails to show the characteris- 
tics of an ordinary chemical change. Something has taken place 
that represents a condition between that of mixture and com- 
pound. Some weak kind of chemical action may have occurred, 
but it seems best at this elementary stage of the study of chem- 
istry to consider solutions as a result of physical changes. 

A property to be especially noted is that solutions are 
homogeneous to the smallest part, and that mechanical methods 
will not separate the substances. Another property of solutions is 
the variation in composition. Salt and water dissolve in all pro- 
portions up to a certain maximum. Alcohol and water mix in all 
proportions without limit either way. A solution containing 5 
per cent salt and one containing 10 per cent salt are not suffi- 
ciently different to be called different substances. 

A solution is defined as a homogeneous mixture of two or more 
substances which cannot be separated by mechanical means and 
ivhose composition may vary. The substance which is thought of as 
going into solution is the solute. The one that does the dissolving 
is called the solvent. 

Solution of gases in liquids. Each state of matter is found 
dissolving in every other state of matter. Most solutions are 
those where a liquid is the solvent. Gases dissolve in liquids — 
some, like oxygen and hydrogen, to a slight extent only; others, 
like ammonia and hydrogen chloride, to a large amount; and still 
others, like carbon dioxide, dissolve more moderately (consult the 



ELEMENTARY CHEMISTRY 43 

table on page 144). The nature of the gas, therefore, determines 
the amount of solubility; so does the nature of the solvent; for 
example, water and alcohol do not dissolve the same amounts of a 
gas. 

Increase of pressure increases the amount of solubility, the 
increase being proportional to tlie pressure. This is the law of 
Henry (page 145). Under high pressure large quantities of gas 
may be held in solution. The excess gas in solution suddenly 
escapes when the pressure is removed, causing effervescence. An 
example is common soda water, or other effervescing drinks, 
which are bottled with the gas dissolved under pressure. Gen- 
erally an increase of temperature decreases the amount of the gas 
which will remain in solution. 

Solids in liquids. Nature of the solid, nature of the solvent, 
and temperature are the factors that influence the amount of 
solid which dissolves. Different substances vary widely in the 
amount which will be dissolved. No solid is entirely insoluble in 
water, but barium sulfate is almost so. Usually, increasing the 
temperature of the solvent will cause it to dissolve more of the 
solid (consult the table on page 147). Calcium sulfate is seen to be 
an exception. 

Molar solutions. Concentration of solutions may be expressed 
by parts per hundred or grams in 100 cc. It is usually more sat- 
isfactory to use a method which shows the formula weight of the 
substance in a certain volume of solution (not of solvent). This 
is called a molar solution. A molar solution is obtained by dis- 
solving the formula weight, or molecular weight, in grams of the 
substance in enough solvent to make 1 liter of solution. 

yThe student must clearly understand that equal volumes of all 
molar solutions contain the same number of molecules. Take two 
substances, one having twice the molecular weight of the other. 
We do not know how many molecules there are in a gram; but, if 
we take 2 gms. of one and 1 gm. of the other, we have taken the 
same number of molecules. An example may make this plainer. 
Assume two piles of shot, one kind weighing 2 gms. each, the 
other weighing 1 gm. Assume that you do not know the number 
of grams in a pound, as above you do not know the number of 
molecules in a gram. If, now, you weigh out 2 lb. of the 2-gm. shot 



44 ELEMENTARY CHEMISTRY 

and 1 pound of the 1-gm. shot, you do not know how many shot 
you have weighed in either case, but it is clear that the number is 
the same in both cases. In like manner, molar, or gram-molecular, 
solutions contain the same number of molecules in equal volumes. 

Saturation. Some liquids dissolve in each other in all pro- 
portions, but with solids (and some liquids too), a point is reached 
where the solvent will dissolve no more. This is a saturated solu- 
tion. When some of the solid remains in contact with the solu- 
tion without change in concentration, the solution is said to be 
saturated. If the solvent evaporates from a solution, the solid 
separates out, usually in crystalline form. If a nearly saturated 
solution is made at high temperature, the solid in excess of the 
quantity required to saturate the solution at ordinary temperature 
usually separates out on cooling. 

In a few cases a condition of supersaturation exists. Sodium 
thiosulfate (Na 2 S 2 3 ), commonly called "hypo," gives a good 
illustration of this. Several grams will dissolve in a few drops of 
water, if heated. The solution may be cooled without crystals 
separating as long as it is undisturbed. Dust particles or a small 
crystal of the "hypo" will cause immediate crystallization. A 
supersaturated solution is one which contains more of the solute in 
solution than the solvent will dissolve at the given temperature. The 
excess of solute, which went into solution at a higher temperature, 
does not separate out on cooling until disturbed in some way. 

Properties of solution. 1. Tendency to uniformity. A solu- 
tion tends to become uniform, the heavy solid rising opposite to 
the force of gravity. This is a similar tendency to diffusion of 
gases. The process is slow unless aided by stirring. Once the 
solution becomes uniform, the uniformity is permanent. 

2. Effect on boiling and freezing points. A solid dissolved in 
a liquid raises the boiling point and lowers the freezing point of 
the solvent. This change in boiling point or freezing point of the 
solvent is proportional to the molecular concentration. (Reread 
paragraph on molar solutions.) Thus, it appears that these 
changes are caused by the number and not by the kind of 
molecules. 

3. Electrolysis. Some solutions conduct the electric current 
and some do not. Salt dissolved in water is a conductor; sugar 



ELEMENTARY CHEMISTRY 45 

dissolved in water is not. A solution which conducts the current 
is called an electrolyte. This conductivity- is always accompanied 
by a decomposition of the substance in solution. This decomposi- 
tion is called electrolysis. The electrodes are the terminals of the 
wires by which the current enters and leaves the solution. The 
positive electrode is called the anode, the negative one is called 
the cathode. 

Freezing points of conductors. A study of solutions shows 
that such substances as sugar, glucose, alcohol, and other organic 
compounds, if soluble in water, form solutions which do not con- 
duct the current, while sodium chloride, copper sulfate, sulfuric 
acid, and other inorganic compounds of the classes of acids, bases, 
and salts do form conducting solutions. It is further noticed that 
those substances whose solutions conduct the current affect the 
freezing and boiling points of the solvent abnormally. 

This will be made clear by a comparison of the behavior of 
solutions in water of sugar, glucose, and salt. Take molar solu- 
tions of each, that is, the molecular weight in grams dissolved in 1 
liter of solution. Such solutions of sugar and glucose will freeze 
at 1.86° below 0° C. All molar solutions which do not conduct 
the electric current give this same lowering of the freezing point. 
A molar solution of common salt, however, freezes at a still lower 
temperature, about 2.79° below zero. Other salts and acids and 
bases give a similar abnormally large lowering of the freezing 
point. If the salt solution is diluted, the actual lowering is less, 
but the actual lowering multiplied by the amount of dilution gives 
a molecular lowering which becomes greater with the greater dilu- 
tion and which approaches a maximum value of two times 1.86°. 
Other compoimds, like sulfuric acid, approach a maximum lower- 
ing of three times 1.86°. 

The theory of ionization. Because organic compounds uni- 
formly lower the freezing point to 1.86° below zero, it was accepted 
that the lowering was caused by the number of molecules present 
in the solution, equal volumes of all molar solutions containing 
equal numbers of molecules. However, compounds giving a still 
lower freezing point are numerous and include such important 
compounds as the acids, bases, and salts, as shown above for 
common salt and sulfuric acid. 



46 ELEMENTARY CHEMISTRY 

Arrhenius proposed to explain these exceptions by assuming 
that such substances in solution have some of their molecules dis- 
sociating into parts, thus forming a greater number of particles to 
produce a greater lowering of the freezing point. Sodium chloride 
molecules break up into two particles. The percentage of dis- 
sociation increases with increased dilution and when all the mole- 
cules are dissociated, we have twice as many particles in the 
solution as original molecules. Therefore, the maximum lowering 
of the freezing point is twice the normal lowering, or twice 1.86°. 
Since sulfuric acid gives a maximum lowering of three times 1.86°, 
each molecule dissociates into three particles in very dilute 
solutions. 

Explanation of conductivity. Until this theory of dissociation 
of molecules in solution was proposed, there was no satisfactory 
explanation of conduction of the electric current by solutions. 
This conduction becomes clear if we assume that the parts into 
which the molecules dissociate are electrically charged, some posi- 
tively and some negatively. If the current is passed through a 
solution of copper chloride, metallic copper is deposited at the 
negative pole and chlorine at the positive pole. It is believed 
that, in the formation of the solution, some of the copper chloride 
molecules dissociate into parts, which at the same time become 
electrically charged. 

These charged parts of the molecule are called ions. The 
copper ions have positive charges, and the chlorine ions negative 
charges. These are in motion in every direction before the current 
is passed. When the current is passed, the positive ions are 
attracted to the negative pole, while the negative ones go to the 
positive pole. Upon arriving at the poles, the ions become dis- 
charged, giving up their charges to the poles. In this way they 
carry the current through the solution. 

It is to be noted that those solutions which have an abnor- 
mally great lowering of the freezing point are the same ones 
which are conductors. The same theory explains both sets of 
facts. 

Statement of the ionic theory. 1. Some molecules of some 
substances when they go into solution dissociate into parts called 
ions. 



ELEMENTARY CHEMISTRY 47 

2. These ions are electrically charged. For that reason their 
properties differ from those of the atom or molecule. 

3. The positive charges always equal the negative charges. 
Some ions carry two or more charges, so the number of positive 
and negative ions is not always equal. Copper chloride (CuCl 2 ) 
dissociates into Cu, carrying two positive charges, and into two 
chlorine ions, each carrying a single negative charge. Positive 
ions are called cations. Negative ions are called anions. 

4. All compounds do not ionize. Those that do ionize do so 
in varying degrees. Hydrogen chloride is ionized by water to a 
large degree, while acetic acid is but slightly ionized, and sugar 
not at all. 

5. Not all solvents produce ionization. Benzene will not 
ionize hydrogen chloride. The same compound is ionized to dif- 
ferent amounts in different solvents. Water is the best ionizing 
solvent; alcohol also produces ionization, but to a much less 
extent than water. Dilution of the solution increases the per- 
centage of ionization. 

Electrolysis. The theory of ionization explains the electro- 
lytic decomposition of substances in solution. This method is used 
in copper, nickel, and silver plating and in electrotyping. A few 
examples will be given. 

With a solution of copper chloride, the copper ions carrying 
two positive charges travel to the cathode when the current is 
turned on. They give up their charge to this pole and deposit on 
it as molecules of copper, since copper has no action on the water 
present. To the anode pole chlorine comes and loses its charge. 
The chlorine escapes as a gas. It has only a slight action with 
the water, if the solution is concentrated. 

Sulfuric acid was used in the so-called electrolysis of water. 
In reality it was the sulfuric acid that was decomposed by the 
current and not the water. That the elements of water were 
liberated was due to a secondary action at the anode. In solu- 
tion sulfuric acid forms two positive hydrogen ions and one sulfate 
ion (SO4) carrying two negative charges. When the current is 
passed, two hydrogen ions go to the cathode, are discharged, and 
are liberated as a molecule of hydrogen. For each molecule of 
hydrogen, a single sulfate ion is discharged at the anode. Here it 



48 ELEMENTARY CHEMISTRY 

acts with water to form sulfuric acid, and one atom of oxygen, 
which of course forms molecules, escapes as a gas. 

Sodium hydroxide is a base. Its molecule ionizes to form 
Na + and OH - ions. At the cathode the sodium is discharged 
when a current of electricity is turned on. Unlike the copper it 
does not deposit, neither does it escape like hydrogen. Sodium is 
a very active metal chemically and reacts with the water of the 
solution to form hydrogen, according to the following equation: 

2Na+2HOH->2NaOH+H 2 

At the anode the OH ions become discharged but do not escape 
as such. They combine to form water and oxygen: 

40H->2H 2 0-f0 2 

Properties of electrolytes. The properties of an electrolyte 
will depend upon both the ions and the molecules present. Ions 
are very much more active than molecules, so they usually are 
responsible for the chemical properties of the solution. 

Silver nitrate solution will precipitate silver chloride from 
aqueous solutions of all chlorides because chlorides form the 
chlorine ion. It will not precipitate silver chloride from a water 
solution of potassium chlorate; potassium chlorate contains chlo- 
rine, but, instead of forming the chlorine ion, it forms the ions K + 
and CIO3"". Neither will silver nitrate precipitate the chlorine 
from organic chlorides which do not ionize. Some ions have a 
characteristic color, as Cu ++ is blue, Fe +++ yellow, and Cr +++ 
green. 

EXERCISES 

1. Define the following terms: solution, solute, solvent, saturated solu- 
tion, supersaturated solution, molar solution, electrolysis, electrolyte, cathode, 
anode, ion, cation, anion. 

2. Discuss factors affecting the solubility of a solid and of a gas in 
liquids. 

3. State the effect of a dissolved solid upon the boiling and freezing 
points of solutions. 

4. Explain why molar solutions contain equal numbers of molecules in 
equal volumes. 

5. What classes of compounds form solutions which are conductors? 

6. State the points in the theory of ionization. 

7. Give some reasons in support of the ionic theory. 

8. How do sugar and salt affect the freezing point of their solutions? 



ELEMENTARY CHEMISTRY 49 

9. Explain the electrolysis of copper chloride (C11O2); of sodium 
hydroxide (NaOH). 

10. Solve problems 8, 9, 10, and 11 on page 158. 

11. Answer questions 2, 3, 4, 5, 6, and 7 on pages 157 and 158. 

12. Answer questions 12, 13, and 14 on page 158. 

13. Why is salt brine used in making artificial ice? 



LESSON IX 

CHLORINE; SODIUM 

CHLORINE 

Assignment: Chapter XIV, McPherson and Henderson 

Occurrence. Chlorine is a very active non-metallic element. 
Because of its great activity it is found in nature only in the com- 
bined state, never as the free element. In this respect it is 
unlike the elements already studied. Sodium chloride is the most 
abundant compound of chlorine, and it is from this substance that 
chlorine is made. 

Preparation. For commercial purposes chlorine is prepared by 
the electrolysis of a concentrated solution of sodium chloride. 
Chlorine escapes at the anode and sodium hydroxide (NaOH) is 
formed at the cathode by the action of sodium with water (pages 
155 and 161). Sodium chloride is cheap, and sodium hydroxide 
has many uses (see pages 175 and 176). The chlorine obtained is 
either compressed in metallic cylinders or absorbed in bleaching 
powder (page 427), from which it can easily be obtained. 

For preparing chlorine in small quantities, as in a laboratory, 
two other methods are used. Both involve the setting free of 
chlorine from hydrochloric acid. The substances used to do this 
are manganese dioxide and potassium permanganate. The reac- 
tions may be explained as oxidizing actions, the manganese dioxide 
(Mn0 2 ) and potassium permanganate (KMn0 4 ) furnishing the 
oxygen to combine with the hydrogen of the acid to form water 
and set free the chlorine. 

In the case of manganese dioxide another explanation is possi- 
ble, as is shown by the equations on page 159. It is frequently 
the practice to make the hydrochloric acid at the time when, and 



50 ELEMENTARY CHEMISTRY 

in the apparatus where, the chlorine is to be prepared. Sodium 
chloride, sulfuric acid, and manganese dioxide are added. The 
following equation represents the final result of the reaction: 

Mn0 2 +2NaCH-2H 2 S04-^Na 2 S04+MnS04+2H20+Cl 2 
Sometimes free oxygen is used to set chlorine free from hydro- 
chloric acid, but the yield is small. 

Properties. Chlorine is a gas, yellowisn-green in color. It 
can be easily liquefied, is 2.5 times as heavy as air, soluble in 
water, has a suffocating odor, and is very irritating upon the nose, 
throat, and lungs (page 162). 

Chemical conduct. Chlorine is a very active element with 
negative, or non-metallic, properties. Like oxygen it combines 
with most of the other elements but it is much more active than 
oxygen at ordinary temperatures. Many of the elements, like 
phosphorus, arsenic, antimony, iron, and sodium, burn in it to 
form chlorides. Even gold and platinum are tarnished by it. 

It will vigorously support the combustion of hydrogen, giving 
a greenish-yellow flame and forming hydrogen chloride (HC1). 
Light will cause a mixture of hydrogen and chlorine in equal 
volumes to explode. Chlorine has such a strong tendency to com- 
bine with hydrogen that it will extract it from compounds con- 
taining it. Sulfur and hydrochloric acid are formed when chlorine 
is passed into a solution of hydrogen sulfide 

C1 2 +H 2 S->2HC14-S 

Turpentine decomposes with the formation of soot when a piece of 
paper, moist with it, is introduced into a jar of chlorine, hydro- 
chloric acid (LIC1) and carbon being formed. Chlorine will set 
oxygen free from water. Sunlight increases this action (page 164). 
Chlorine solution in water is, therefore, a good oxidizing agent. 

Chlorine is a strong bleaching agent, but water is necessary to 
the action as experiment shows that dry chlorine will not bleach 
pieces of colored calico which are quickly bleached if wet (page 
165). It would seem, therefore, that it is the oxygen set free 
from the water that is the immediate agent in bleaching the 
cloth. Only those colored substances can be bleached whose col- 
oring matter can be oxidized to colorless compounds. Printer's 
ink, which is made of carbon, cannot be bleached (page 165). 



ELEMENTARY CHEMISTRY 51 

Uses. Besides being used largely as a bleaching agent, chlo- 
rine is used as a germicide or disinfectant in the form of either the 
free element or compounds from which it is readily set free, like 
bleaching powder. Cotton cloth is now bleached by the use of 
chlorine. It is used to kill microorganisms in drinking water, as 
in city water supplies, and to purify water for the army. Chlo- 
rine was the first of the cloud gases used in the Great War, and 
later many of the gas bombs contained it or some of its com- 
pounds. 

Nascent state. The oxygen set free by the action of chlorine 
on water is much more active than ordinary oxygen. It bleaches 
colored cloth which would not be bleached by pure oxygen or by 
air. The difference is due to the fact that the bleaching is done 
just at the time when the oxygen is set free from combination 
with water and before it has combined with itself to form mole- 
cules of oxygen. It is at this time in the atomic state, and none 
of its energy is used up in holding its atoms in a molecule. In 
this atomic state it is much more active. This state is called the 
nascent state. It is the state existing at the time of its liberation 
from compounds. The root of the word "nascent" means to be 
born. 

HYDROGEN CHLORIDE 

Preparation. The compound of hydrogen and chlorine may 
be made by the direct union of chlorine and hydrogen as has 
been shown. Hydrogen chloride can be more conveniently made 
by another method. Common salt (NaCl) is treated with sulfuric 
acid (H 2 S0 4 ). Sodium sulfate (Na 2 S0 4 ) and hydrogen chloride are 
formed: 

2NaCl+H 2 S0 4 -+Na 2 S0 4 +2HCl 

For commercial purposes it can be prepared by using the same 
reaction (page 167). It is also prepared commercially from sodium 
chloride (NaCl) and sodium hydrogen sulfate (NaHS0 4 ) (page 
168): 

NaCl+NaHS(V->Na 2 S0 4 +HCl 

The hydrogen chloride escapes as a gas and is dissolved in water. 
This solution is the hydrochloric acid of commerce. When pure it 



52 ELEMENTARY CHEMISTRY 

is colorless; but frequently it is colored yellow by impurities. The 
impure acid is called muriatic acid, meaning the acid of sea salt. 

Properties. Its solubility in water, irritating odor, and gase- 
ous state are its important physical properties. One volume of 
water dissolves 506 volumes of hydrogen chloride gas. Other 
physical properties will be found stated on page 169. 

Composition. By the electrolysis of a solution of hydrogen 
chloride, or by combining the two gases, it can be shown that 1 
volume of hydrogen and 1 volume of chlorine combine to form 2 vol- 
umes of hydrogen chloride. Like the volume composition of water 
this fact must be remembered for future use in selecting the 
standard of molecular weights. By weight, hydrogen chloride 
contains 35.18 parts of chlorine and 1 part of hydrogen to every 
36.18 parts of the gas. 

Hydrochloric acid. Hydrogen chloride as a gas is little used 
and is inactive. Its solution in water, hydrochloric acid, is very 
active because of the formation of hydrogen and chlorine ions. It 
is one of the common acids (page 169). As such it has a sour 
taste, turns litmus red, and is acted upon by certain metals with 
the replacement of its hydrogen by the metal. Iron, zinc, and 
sodium so act. The equations for these actions are found on page 
171. It acts with bases, like sodium hydroxide (NaOH), to form 
water and a salt: 

HCl+NaOH->H 2 0+NaCl 

Hydrochloric acid under certain conditions is a reducing 
agent — not so active a reducing agent as some other compounds, 
but it will give up its hydrogen to compounds like nitric acid and 
potassium permanganate. The equations for these actions will be 
taken up when these compounds are studied. Hydrochloric acid 
is a very strong acid. What is meant by the strength of an acid 
will be learned in the next lesson. 

SODIUM 

Assignment: Chapter XV, McPherson and Henderson 

Metals. Elements are divided into two great classes, metals 
and non-metals. Chemically the two classes have opposite proper- 
ties, and they have wide differences in physical properties. The 



ELEMENTARY CHEMISTRY 53 

metals are solids, except mercury which is a liquid, they have a 
metallic luster, and they conduct heat and electricity. The chem- 
ist sometimes distinguishes between the chemical properties of the 
metals and the non-metals by describing the metals as positive 
and the non-metals as negative. A metal is an element which may 
become a positive ion upon solutions being made of its compounds. 
In this sense hydrogen is a metal, though it has the physical 
properties of a non-metal. 

Occurrence of sodium. Sodium is a widely distributed ele- 
ment, but because of its great chemical activity, especially toward 
oxygen and water, it is not found in the free state, but only in its 
compounds. Its compounds are all soluble in water, so they are 
largely found in sea water, as for example, common salt, or 
sodium chloride. This compound is the one used for preparing 
the metal. Chili saltpeter (NaN0 3 ), borax (Na 2 B 4 7 ), and sodium 
carbonate (Na 2 C0 3 ) are also found in nature. 

Preparation. Sodium was first made by Davy by the elec- 
trolysis of fused sodium hydroxide. Prior to this work sodium 
hydroxide was thought to be the element. The commercial 
preparation of sodium today utilizes the same method, in prin- 
ciple. A different source of electrical energy is used; electrical 
energy is now comparatively cheap, as it is generated by the 
dynamo, and particularly so if the dynamo is run by water 
power, as at Niagara Falls. Davy had to obtain his energy from 
a large battery of wet cells, making the method very expensive. 

Sodium cannot be obtained by the electrolysis of its com- 
pounds in solution, because of the action of the element with 
water. The process in use at Niagara Falls was devised by 
Castner (page 174) and is as follows: Sodium hydroxide is 
melted and then becomes a conductor of the current, sodium 
being deposited at the cathode and oxygen escaping from the 
anode. 

Prior to the generation of cheap electrical energy and its 
application to chemical industries, sodium was made by reducing 
sodium oxide with carbon. To bring about this reaction sodium 
carbonate was heated with carbon. This equation shows the 
reaction: 

Na 2 C0 3 +2C->2Na+3CO 



54 ELEMENTARY CHEMISTRY 

In reality this occurs in three steps: 

Na 2 C0 3 heated->Na 2 0+C0 2 

Na 2 0+C->2Na+CO 

C0 2 +C->2CO 

Properties and uses. Sodium is so active that it has no uses 
as an element. It is one of the strongest, or most active, metals 
and forms with water one of the strongest bases, sodium hydroxide 
(NaOH). With air the action is so great that it must be kept 
under kerosene (page 175). Sodium compounds impart a deep 
yellow color to a colorless or bluish flame, if a platinum wire is 
dipped into a solution containing them and then held in the 
flame. This serves as a test for the metal in its compounds. 

Preparation of sodium hydroxide. As shown in the preceding- 
paragraph, sodium hydroxide is obtained by the action of sodium 
on water. In its commercial preparation use is made of this 
reaction. The method consists of the electrolysis of a sodium 
chloride solution. The sodium discharged at the cathode reacts 
with the water to form sodium hydroxide. (See the description of 
the Townsend cell on page 176.) An older method consists of 
using moist lime (Ca(OH) 2 ) and sodium carbonate (Na 2 C0 3 ). 
The equation is 

Na 2 C0 3 +Ca(OH) 2 ->CaC0 3 +2NaOH 

A variation of this method applied to potassium compounds 
formerly was a familiar farm-house operation. In the earlier days 
nearly every household made its own soap. Lye, either sodium 
hydroxide or potassium hydroxide, is necessary to act upon the 
fat and convert it into soap (see "Soap" in Lesson XVIII). To 
obtain this lye, wood ashes (preferably hardwood ashes) were 
treated with moist lime in a so-called "ash hopper," which was an 
inverted pyramid of rough boards lined with straw. Into this the 
mixed ashes and lime were put. Rain or added water promoted 
the reaction, the lye solution being collected as drippings at the 
bottom. Wood ashes contain potassium carbonate (K 2 C0 3 ), so 
the lye obtained was an impure solution of potassium hydroxide. 
It made a soft soap, while sodium hydroxide makes a hard soap. 

Properties and uses. Sodium hydroxide is a white solid, very 
active and, therefore, hard to keep pure. It is a very strong base. 



ELEMENTARY CHEMISTRY 55 

It has a corrosive action on animal and. vegetable matter and is, 
therefore, called caustic soda. It is used where a strong base is 
necessary, in making soap, paper and dyes, in bleaching, and in 
refining kerosene oil. 

EXERCISES 

1. Describe the methods for preparing chlorine. 

2. Give the physical properties of chlorine. 

3. Discuss completely the chemical conduct of chlorine. 

4. Give the uses of chlorine. 

5. Explain nascent state. Give an example. 

6. How is hydrogen chloride prepared? 

7. Give the properties of hydrogen chloride. 

8. What is the volume and the weight composition of hydrogen chlo- 
ride? How are these facts determined? 

9. Discuss the properties of hydrochloric acid. 

10. Answer questions 1, 2, 4, 7, and 9, page 171. 

11. Solve problems 3, 5, 6, 8, and 10, page 171. 

12. What is a metal? 

13. Give the methods for preparing sodium. 

14. Describe the Castner process for making sodium. 

15. Give the properties of sodium. 

16. Give the chemical name and the formula for caustic soda. 

17. How is sodium hydroxide made? 

18. Discuss the properties and uses of sodium hydroxide. 

19. Solve problems 4, 5, 6, 7, and 9, page 178. 

20. Answer questions 8 and 10, page 178. 

21. How would you test for sodium compounds? 

LESSON X 
ACIDS, BASES, AND SALTS 
Assignment: Chapter XVI, McPherson and Henderson 

Introduction. In the study of that branch of chemistry 
known as inorganic chemistry (it is with this branch that this 
course principally deals), the great majority of the compounds 
that can be studied belong to one of the groups known as acids, 
bases, and salts. Another important group is that of the oxides. 
Lavoisier considered oxygen to be a necessary constituent of an 
acid, as is evidenced by his naming the element "oxygen" (acid 
producer) and by his defining an acid as a compound of oxygen. 
We now recognize that in the case of such acids as hydrochloric 
acid oxygen is not present. 



56 ELEMENTARY CHEMISTRY 

Oxides are closely related to both the bases and the acids/ 
There are two kinds of oxides, oxides of metals and oxides of non- 
metals. Metallic oxides are related to bases as is shown by the 
following equation: 

Na 2 0+H 2 0->2NaOH 

Non-metallic oxides are related to acids in a similar way: 

S0 3 +H 2 0->H 2 S0 4 
Sodium hydroxide (NaOH) and sulfuric acid (H 2 S0 4 ) are a base 
and an acid respectively. 

The common acids. The word "acid" means a sour sub- 
stance. The earliest acid known was acetic acid, which is the 
sour constituent in vinegar. The name "acetic" comes from the 
same word root, meaning sour. Acetic acid is formed from fruit 
juices by the action of certain bacteria, hence its early recognition. 
It is one of a large number of organic acids. 

Just now we are more concerned with the inorganic acids, or 
mineral acids, as they are sometimes called. Common among 
these are hydrochloric (HC1), sulfuric (H 2 S0 4 ), and nitric (HN0 3 ) 
acids. These are usually used in water solution, and their acid 
properties are active only in solution. Generally, it is not neces- 
sary to make a distinction between the pure substance and its 
solution. Both are called acids. 

Characteristics of acids. All acids have certain common 
properties. They all have a sour taste; they act upon certain 
organic substances to change the color (litmus is changed to a red 
color by acids); their solutions are conductors of the electric cur- 
rent, hydrogen being set free at the cathode. Certain metals 
have the power to liberate hydrogen from acids. Thus 

Zn-{-2HCl->ZnCl 2 +H 2 

and 

Fe-hH 2 S0 4 ->FeS0 4 +H 2 

Copper and some other metals will not liberate hydrogen. All 
acids will react with metallic hydroxides (bases) to form water and 
a salt, as shown by these equations: 

NaOH+HN0 3 -»HOH-fNaN0 3 
Fe(OH) 2 +H 2 S0 4 -*2HOH+FeS0 4 
Cu(OH 2 ) +2HCl->2HOH+CuCl a 



ELEMENTARY CHEMISTRY 57 

Properties of compounds are the result of the composition of 
the compounds. Since all acids contain hydrogen, it may be con- 
sidered that these common properties are due to hydrogen. The 
properties are active only when the acid molecule is dissociated in 
solution; we may then assume further that the characteristic 
properties of acids are due to the presence of hydrogen ions (H + ). 
An acid may be defined as a compound which, when in solution in a 
dissociating solvent, produces hydrogen ions. 

Common bases. The common bases are frequently called 
alkalies. The word "alkali" originally meant ashes and was first 
applied to the ashes of sea plants. It is now applied to the 
hydroxides and carbonates of sodium, potassium, and sometimes 
to those of magnesium and calcium. The term "vegetable alkali" 
referred to the ashes of plants, which usually contained potassium 
compounds. The sodium compounds, obtained from rock salt, 
were called mineral alkalies. 

The carbonates were formed in the ashes and were the first 
alkalies observed. As shown in Lesson IX, the carbonates treated 
with lime, that is, burnt limestone, are changed to the hydroxides. 
The hydroxides are much more active than the carbonates and 
have a stronger alkali action. The carbonates are called mild 
alkalies, while the hydroxides are called caustic alkalies. The 
hydroxides of sodium and potassium are commonly found in the 
grocery store under the names "soda lye" and "potash lye." 

Characteristics of bases. Bases have a bitter taste, if soluble; 
they turn litmus blue; their solutions are conductors, a metal 
being discharged at the cathode and oxygen at the anode; they 
react with acids to form water and a salt. It is found that they 
all contain a metal and oxygen and hydrogen in the group known 
as hydroxyl (OH). They dissociate in solution to form a cation of 
the metal and an anion of the hydroxyl (OH). A base may be 
defined as a compound which, when in solution in a dissociating 
solvent, produces hydroxyl ions. The properties of bases, then, are 
due to the presence of these ions. 

Undissociated acids and bases. These compounds in a per- 
fectly dry state do not have the properties of acids and bases. 
Dry hydrogen chloride is not acted upon by zinc to form hydro- 
gen. Hydrogen chloride, or hydrochloric-acid gas, dissolved in 



58 ELEMENTARY CHEMISTRY 

benzene does not act with zinc to form hydrogen. Such a solu- 
tion is not a conductor of the electric current. Benzene does not 
dissociate the hydrogen chloride molecules into ions and, hence, 
does not form a solution with acid properties. Similar statements 
apply to bases. 

Neutralization. As shown under the properties of acids and 
bases, they react with each other to form water and a compound 
known as a salt. In this action the acid properties and basic 
properties are lost. If proper proportions of the acid and the base 
are used, the resulting mixture has neither a sour nor a bitter 
taste, turns litmus neither red nor blue. For this reason such a 
chemical reaction is called neutralization. 

What really happens is that the hydrogen ion of the acid 
reacts with the hydroxyl ion of the base to form the undissociated 
molecule of water. The metallic ion of the base and the non- 
metallic ion of the acid remain ions, if the solution is dilute, or 
unite to form molecules, if the solution is more concentrated. 
These equations represent some neutralizations: 

2KOH+H 2 S0 4 ->2HOH+K 2 S0 4 
Ca(OH) 2 +2HCl->2HOH-f-CaCI 2 

Neutralization may be defined as the action of a base and an acid in 
which the cation of the acid unites with the anion of the base to form 
an undissociated molecule of water. The formation of the salt is a 
secondary step. Evaporation of the water permits the ions of the 
salt to come together to form molecules. 

Neutralization, a definite act. Neutralization is a definite 
chemical action between base and acid compounds. Housewives 
are sometimes surprised that sugar does not sweeten a sour sub- 
stance, while soda will. When milk sours, a definite acid known as 
lactic acid is formed. Sugar will not correct the unpleasant sour 
properties this gives to the milk. Soda will, for soda is a mild 
alkali, that is, it has basic properties; it corrects the sour taste of 
the milk by converting the lactic acid into water and a salt. 
Sugar, on the other hand, has no chemical action with the lactic 
acid. It leaves the lactic acid, and consequently the sour taste, in 
the milk. Large quantities of sugar may mask the sour taste, but 
not remove it. 



ELEMENTARY CHEMISTRY 59 

The use of soda and sour milk in baking is an example of a 
similar nature. Here the purpose is to form carbon dioxide to 
raise the dough. The action must be between definite quantities 
of soda and sour milk. If not enough soda is used, the sour 
taste remains; if too much soda is used, the result is a bitter 
taste and a yellow color. The experiment described on pages 183 
and 184 with solutions of acid and base of known concentration 
shows neutralization to be a definite chemical act. If 5 cc. of the 
acid solution requires 8 cc. of the base solution to neutralize it, 
then 10 cc. of the acid will require 16 cc. of the base, 20 cc. of the 
acid will require 32 cc. of the base, and 30 cc. of the acid will 
require 48 cc. of the base to neutralize them. 

Heat of neutralization. That the real chemical act in neutral- 
ization is the formation of water, and nothing but water, is shown 
by the examples given under this heading on page 184. 

Salts. Sodium chloride is the commonest of all salts and 
gives the name "salt" to the group. Salts have a peculiar taste, 
which may be described as salty, but that does not mean the 
same as the taste of common salt. As a rule they have no action 
on litmus; their solutions are conductors; they are composed of a 
metal and a non-metal or a group of non-metal atoms. The fol- 
lowing equations will show the formation of some salts: 

Mg(OH) 2 +2HCl->2HOH+MgCl 2 

Zn(OH) 2 +2HN0 3 ->2HOH+Zn(N0 3 )2 

2Al(OH) 3 +3H 2 S0 4 ^6HOH+Al 2 (S0 4 ) 3 

Salts are also formed when a metal replaces hydrogen: 

2A1+6HC1->2A1C1 3 +3H 2 
Mg+H 2 S0 4 ->MgS0 4 +H 2 

The relation of salts to acids and bases is close. The salts 
contain the metal of the base and all of the acid except the hydro- 
gen. A salt may be defined as a compound which is formed by the 
union of the cation of the base with the anion of the acid. Solutions 
of salts have no common ion. 

Normal salts, acid salts, basic salts. A normal salt is one in 
which all of the hydrogen of the acid has been replaced by metal, and 
all of the hydroxyl of the base has been replaced by non-metallic ions, 
examples: NaCl, Na 2 S0 4 , Mg(N0 3 ) 2 , CaCl 2 . Evidently an acid 



60 ELEMENTARY CHEMISTRY 

that contains only one hydrogen atom can form only a normal 
salt. One like sulfuric acid (H 2 S0 4 ) presents two possibilities. 
One or both of the hydrogen atoms may be replaced as follows: 

NaOH+H 2 S0 4 ->HOH+NaHS0 4 
2NaOH+H 2 S0 4 ->2HOH+Na 2 S0 4 

A compound like sodium hydrogen sulfate (NaHS0 4 ) is called 
an acid salt. An acid salt may be defined as a salt which still con- 
tains a part of the hydrogen of the acid from which it is derived. 
Acid salts do not necessarily turn litmus red. Other reactions 
may cause an acid salt like baking soda (NaHC0 3 ) to turn litmus 
blue. 

A base which contains more than one hydroxyl group may 
form basic salts in like manner: 

HCi+Fe(OH) 3 ->Fe(OH) 2 Cl+HOH 

A basic salt is a salt which contains a part of the hydroxyl of the 
base from which it is derived. In solution acid salts and basic salts 
form ions characteristic of the acid or base as well as of the salt. 

Strength of acids and bases. The strength of acids and bases 
must not be confused with the concentration of their solutions. 
Acid and basic properties are due to the presence of hydrogen and 
hydroxyl ions. That one will exhibit its characteristic properties 
to the greatest degree which dissociates the greatest percentage of 
its molecules into ions. Dissociation depends upon the nature of 
the compound and upon the concentration of the solution. In 
order to compare the strength of acids, their solutions must be 
taken of such concentration that they contain equivalent weights 
of hydrogen; for example, if a molar solution of hydrochloric acid 
is used, a half-molar solution of sulfuric acid must be used. 

Conductivity measurements of such solutions show that the 
three common mineral acids are quite strong. In the order of 
strength, they are hydrochloric, nitric, and sulfuric. Other acids 
vary considerably in strength. Acetic acid is a weak acid. Bases 
in like manner vary considerably in strength. Potassium hydrox- 
ide is the strongest of the common bases, and sodium hydroxide 
comes next. Dilution of the solution increases the percentage of 
molecules which dissociate, but of course the actual number of 



ELEMENTARY CHEMISTRY 61 

ions in a given volume is less as dilution is greater. Some sub- 
stances dissociate in more than one way; take, for example, sul- 
furic acid, which forms the ions H + and HS0 4 ~ in concentrated 
solution, and H + , H+, S0 4 ~ ~ in dilute solution. 

Radicals. Groups of atoms like OH, S0 4 , N0 3 , NH 4 , and 
C0 3 , which are found in a number of related compounds and pass 
through chemical changes as a unit, are called radicals. The 
radicals of acids, bases, and salts become the ions upon dissocia- 
tion. A radical is defined as a group of atoms acting as a unit in 
chemical change. 

Nomenclature. The reader should refer to pages 189 and 190 
for the methods used in naming acids, bases, and salts. It only 
needs to be added here that while all bases are hydroxides, all 
hydroxides are not bases. In fact all the oxygen acids are con- 
sidered to be hydroxides. This is shown if we write the formulas 
as follows: 

H 2 S0 4 =(HO) 2 S0 2 
HN0 3 =HON0 2 

Electrochemical series. Reference has been made more than 
once to the fact that some metals, like zinc and iron, will replace 
hydrogen from acids, while others, like copper and silver, will not. 
Also, zinc will replace copper, lead, iron, and silver in solutions of 
their salts: 

Zn+Pb(N0 3 ) 2 ^Pb+Zn(N0 3 ) 2 

Copper will replace silver and mercury, but will have no effect 
upon solutions of salts of lead, iron, tin, and zinc. 

A study of all the metals and hydrogen shows that they can 
be arranged in a series in the order of their ability to replace from 
solutions of their salts metals farther down the series. Study the 
table on page 191. This table of the metals indicates the order of 
their chemical activity. Those metals near the top form strong 
bases. Those above hydrogen will replace hydrogen from acids, 
while those below it will not. 

EXERCISES 

1. Give the characteristic properties of acids, bases, and salts. 

2. Explain how acids and bases are related to oxides. 

3. Define acid, base, salt, neutralization. 



62 ELEMENTARY CHEMISTRY 

4. What is meant by a normal salt, acid salt, basic sale? Give 
examples. 

5. What is meant by the strength of an acid? What is a radical? 

6. What does the heat of neutralization show? 

7. Give the methods of naming acids, bases, and salts. 

8. What is meant by the electrochemical series? Name ten metals in 
the correct order in this series. 

9. What does the term "alkali" mean? What is soda lye; potash lye? 
Distinguish between caustic and mild alkali. 

10. Give examples to show that a dry acid does not have acid properties. 

11. Give some examples to show that neutralization is a definite chemi- 
cal action. 

12. Answer questions 3, 7, 8, 9, and 10, page 192. 

13. Solve problems 4, 5, and 6, page 192. 

SEND EXERCISES FOR LESSONS VI=X TO THE SCHOOL 



LESSON XI 

VALENCE; EQUATIONS 
Assignment: Chapter XVII, McPherson and Henderson 

Definition of valence. A review of the formulas of some of 
the compounds studied in previous lessons will reveal the fact that 
all atoms do not hold in combination the same number of other 
atoms. One oxygen atom combines with two of sodium, but with 
only one of calcium, while two oxygen atoms combine with one of 
carbon. One atom each of chlorine, oxygen, nitrogen, and carbon 
combine respectively with one, two, three, and four of hydrogen. 

Since one atom of hydrogen is never found in combination 
with more than one atom of some other kind, the hydrogen atom 
is taken as the basis for comparing the combining power of other 
atoms. This combining power, or combining value, is called 
valence. Valence may be defined as the property of an element 
ivhich determines the number of hydrogen atoms its element will hold 
in combination. Elements which combine with one hydrogen atom 
are univalent, as shown in the formulas HO, HBr. Bivalent ele- 
ments are those that combine with two hydrogen atoms, as in 
formulas H 2 0, H 2 S. The formulas NH 3 , PH 3 , AsH 3 show trivalent 
elements. No element is known which has a valence higher than 
eight. Carbon is quadrivalent, as in CH 4 . 



ELEMENTARY CHEMISTRY 63 

Construction of formulas. Elements which do not form com- 
pounds with hydrogen may be compared with chlorine, which is 
univalent, or with oxygen, which is bivalent, as in H 2 0. The 
oxides of many elements, then, are used to determine the valence 
of the element. If the formula of the oxide is found by experi- 
ment to be PbO, then the valence of lead is two. If we find it to 
be Pb02, the valence of lead is four. Elements which have the 
same valence, or valences which are multiples of each other, com- 
bine in a simple way. One atom of one element combines with 
one atom of the other, or one atom of one element combines with 
two atoms of the other, as in PbO and Pb0 2 above. 

But, if the valences are not multiples, but are two and three, 
for example, the formula for the compound must be constructed on 
the basis of the least common multiple. Phosphorus has a valence 
of three (PH 3 ) and oxygen two (H 2 0). Each element in phos- 
phorus oxide must be used a sufficient number of times to make 
the total valence of all atoms of each kind equal to the least com- 
mon multiple of the two valences. In this case the least common 
multiple is six. Therefore, the formula for phosphorus oxide is 
P 2 3 . The rule to follow is: Take the least common multiple of the 
valences, divide the least common multiple by the valence, and the 
quotient is the number of times that atom is to be used in the formula. 

Finding valence of elements. The valence of an element 
should not be found by referring to some table of valences, but by 
looking up some formula containing the element in the textbook, 
consulting the index to find the page where that element is 
described. 

Suppose you wish to write the formula for bismuth sulfide. 
The index will refer you to page 362. Here you will find the for- 
mulas BiCl 3 and Bi 2 3 , which show bismuth to be trivalent. 
Knowing the formula of hydrogen sulfide to be H 2 S, or finding it 
on page 233 if it is not known, sulfur is shown to be bivalent. 
The least common multiple would be six, and two atoms of bis- 
muth and three of sulfur are necessary to form the compound 
bismuth sulfide, the formula for which is Bi 2 S 3 . 

If the formula of a compound is known to be As 2 S 5 , the 
valences are found by the opposite procedure. The least common 
multiple of the number of both atoms is ten. The valence of 



64 ELEMENTARY CHEMISTRY 

each element will be ten divided by the number of atoms of that 
element in the formula, or ten divided by two equals a valence of 
five for arsenic and ten divided by five equals a valence of two 
for sulfur. 

Valence of radicals. When a compound consists of more than 
two elements, the above method for determining the valence of 
each element is not applicable. A more extended study of the 
chemical reactions is necessary to determine the manner in which 
the several elements are united to each other. Such a study is 
not possible here. Until such experimental evidence is at hand, 
we could make several guesses as to the valence of sulfur in sul- 
furic acid. See the structural formulas on page 196. 

For our present purposes it is not necessary to know the 
valence of all the individual atoms. We are going to deal prin- 
cipally with acids, bases, and salts, which dissociate into cations 
and anions. These ions are of two kinds only, positive and nega- 
tive. It is the valence of these parts, or radicals, of the molecule 
that we must know. The method described above for a compound 
of two elements like As 2 S 5 applies just as well to one composed of 
a metal and a negative radical, like Fe 2 (S0 4 )3. From this formula 
we could learn that iron is trivalent and the sulfate radical 
bivalent. 

Finding barium to be bivalent from the formula BaCl 2 and 
the phosphate ion (P0 4 ) to be trivalent from the formula H 3 P0 4 , 
we take the least common multiple, which is six, and dividing six 
by the valences, we get three barium atoms and two phosphate 
radicals to provide the total of six positive and six negative 
valences. The formula would be Ba 3 (P0 4 )2. It is clear, then, that 
valence corresponds in number and kind to the charges which the ions 
carry when we are dealing with electrolytes. 

Variable valence. The same element does not always have 
the same valence. When two elements form more than one com- 
pound, one or the other frequently has different valences. Iron 
forms two chlorides, having the formulas FeCl 2 and FeCl 3 . 
Assuming chlorine to have a valence of one in each case, iron is 
two in the former and three in the latter. Carbon has a valence 
of four in C0 2 and of two in CO, if oxygen is two in both 
cases. 



ELEMENTARY CHEMISTRY 65 

This variation in valence need not confuse us. Ordinary 
chemical changes take place without change of valence. Only 
those changes that are oxidation and reduction reactions produce 
changes in valence. These cases will be pointed out as they are 
met with. Most elements have a dominant valence that more 
frequently occurs than the others. Thus carbon is usually quadri- 
valent, lead bivalent, and phosphorus pentavalent, rather than 
bivalent, quadrivalent, and trivalent respectively. 

Equations. The valence of the elements can now be put to 
use to aid us in writing equations. As has been shown in an 
earlier lesson, the chemical equation is merely a representation of 
the facts of the reaction. These facts must always be determined 
either by laboratory experiment or by study of the results of the 
experiments of others. 

Suppose we are required to write the equation representing 
the reaction in which aluminium hydroxide dissolves in sulfuric 
acid. A complete and original investigation of this reaction 
would involve an analysis of aluminium hydroxide and sulfuric 
acid to find formulas to represent their composition. It would be 
found that the formulas Al(OH) 3 and H 2 S0 4 represent these com- 
pounds. Then, the substances formed would have to be studied 
to find out their composition. It could be shown that water and 
aluminium sulfate are formed and that their formulas are H 2 
and A1 2 (S0 4 ) 3 . 

Some of these steps can be predicted, instead of being 
actually determined by experiment, if we apply the knowledge we 
have gained of the principles of valence and the properties of 
compounds. It will be recognized that aluminium hydroxide, 
being a metallic hydroxide, is a base. Sulfuric acid will be recog- 
nized as an acid. From a knowledge of their properties, we know 
that they will act to form water and a salt, as all acids and bases 
will. If we are familiar with the formulas for other compounds of 
aluminium, we will learn that its atom has a valence of three. 
The formula for aluminium hydroxide would therefore be Al(OH) 3 , 
since the OH radical is always univalent. As sulfuric acid has the 
formula H 2 S0 4 , the S0 4 radical is bivalent. A temporary equation 
can now be constructed: 

A1(0H) 3 +H 2 S0 4 -*H0H+A1S0 4 



66 ELEMENTARY CHEMISTRY 

Now apply the valences to obtain the correct formulas on the 
right-hand side of the arrow. The valences give a least common 
multiple of six. The trivalent atom Al must, therefore, be used 
two times, and the bivalent radical S0 4 three times in the formula 
for aluminium sulfate. H 2 or HOH is, of course, the correct 
formula for water. The equation now stands as 

Al(OH) 3 +H 2 S0 4 ->HOH-j-Al 2 (S0 4 ) 3 

It remains to use as many of each of the above molecules as 
are necessary to give an equal number of all atoms on both sides 
of the equation. That is to say, the equation must be balanced. 
In balancing an equation the student must be careful not to 
change the formulas for any of the compounds. Those formulas 
represent the composition of the compounds as nature made them. 
The student must not take upon himself the supernatural power 
to change them. All he can do is to take more of the molecules. 
This is done by placing the proper coefficients in front of the 
formulas. 

It is best to start at the left-hand side of the second tem- 
porary equation, as written above, and, taking each atom or 
radical in turn, count them on each side. Thus, we have one 
Al on the left and two on the right. Therefore, take two mole- 
cules of Al(OH) 3 . This gives us six OH radicals, and six molecules 
of water will be necessary on the right of the equation, since each 
HOH molecule contains one OH radical. Now, we have six 
hydrogen atoms, which will have to be obtained from three 
molecules of H 2 S0 4 . Putting three in front of H 2 S0 4 makes three 
S0 4 radicals, which we find already present in the one molecule of 
A1 2 (S0 4 ) 3 on the right of our equation. All parts having been 
tested and corrected, the complete equation is 

2Al(OH) 3 +3H 2 S0 4 ^6HOH+Al 2 (S0 4 ) 3 

All double decomposition equations can be written by using 
the above method. When the products are not water and a salt, 
it must be known what they are. If a solid precipitates or a gas 
is formed, these substances can be recognized by tests. In the 
questions following this lesson, any reactions in which the products 
are not water and a salt will be described sufficiently so that the 
products will be known. 



ELEMENTARY CHEMISTRY 67 

EXERCISES* 

1. What is valence? 

2. Give four kinds of valence and the formula of a compound to 
illustrate each. 

3. Describe a method for finding the formula of a compound when the 
valences are known. Give an example. 

4. How can you find the valence if you know the formula of a com- 
pound of the element? Give an example. 

5. What is the connection between the valence of radicals and the 
electrical charges or ions? 

6. Describe the steps in writing a double decomposition equation. 
Give an example. 

7. Give complete answers to all the exercises on page 199. 

8. Complete the following equations: 

KOH+HN0 3 -+? 

KOH+H 2 C0 3 -»? 

KOH+H 3 P0 3 -»? 

Ba(OH) 2 +HN0 3 -y? 

Ba(OH) 2 +H 2 C0 3 ->? 

Ba(OH) 2 +H 3 P0 3 ->? 

Cr(OH) 3 +HN0 3 ->? 

Cr(OH) 3 +H 2 S0 4 -*? 

Cr(OH) 3 +H 3 P0 3 -V? 

9. Lead is sometimes quadrivalent. Write the formula for its hydrox- 
ide. Write the equation for the action of lead hydroxide with phosphoric acid 
(H 3 P0 4 ). 

10. From question 8 find the valence of chromium. Construct the 
formula for chromium sulfate. From question 2, page 199, find the valence of 
calcium. Chromium sulfate and calcium hydroxide precipitate chromium 
hydroxide. Write the equation. 

LESSON XII 
COMPOUNDS OF NITROGEN 

Assignment: Chapter XVIII, McPherson and Henderson 

Reference was made in Lesson VII to the occurrence of com- 
pounds of nitrogen. Those parts of Lesson VII dealing with the 
occurrence and properties of nitrogen should be reviewed at this 
time. The chemical conduct of nitrogen compounds depends to a 
great extent upon the chemical properties of nitrogen. It was 
learned that nitrogen is a very inactive element. This may be 

*Prepare this set of Exercises and hold it until those for Lessons XII, XIII, XIV and 
XV are also prepared and then send all five sets to the School, 



68 ELEMENTARY CHEMISTRY 

due to the great stability of the molecule of nitrogen, which has 
the formula N 2 , as we learn by the methods used in Lesson XVI. 

Nitrogen has a tendency to leave its compounds to form the 
more stable nitrogen molecule, with the liberation of energy. 
Therefore, nitrogen compounds show a tendency to be unstable. 
Explosives are compounds of nitrogen which form a large volume 
of gas at the time of their decomposition. Compounds of nitrogen 
are numerous, but only those with hydrogen and oxygen need be 
studied here. 

Ammonia. Three compounds of nitrogen and hydrogen are 
known. Ammonia (NH 3 ) is the only one of sufficient importance 
to require discussion. It is formed by the decay or fermentation 
of organic matter containing nitrogen. For this reason the odor 
of ammonia is noticeable in stables, especially those occupied by 
horses, when the manure heats or ferments. It is sometimes 
called spirits of hartshorn, because it was first prepared by heat- 
ing the horns of animals, such as deer or hart. 

Commercial preparation. Practically all the ammonia used in 
this country is made by the destructive distillation of soft coal. 
The ammonia may be looked upon as a by-product of the process, 
but in reality there are four important products, coke, coal tar, 
ammonia, and coal gas. Soft coal is heated in retorts from which 
air is excluded, and a great variety of substances are obtained 
which belong in the four groups above named. The process is 
described on page 306. 

The gases formed are passed through the liquids in the 
hydraulic main, where the ammonia and water form the 
ammoniacal liquor. Some more ammonia is dissolved by water 
in the scrubber. The ammoniacal liquor is treated with lime, 
and the pure ammonia distills over. It is dissolved in water or 
in dilute acid. Sulfuric acid is frequently used, as it is cheap and 
the ammonium sulfate formed can be used as a fertilizer. 

Recently a process has been developed for making ammonia 
from the elements nitrogen and hydrogen. The reaction is an old 
one. The difficulty with it was that the ammonia decomposed 
almost as fast as it was formed. Haber succeeded in finding con- 
ditions under which this decomposition is retarded. The best 
yield is found when the mixture of gases is heated to 500° C. 



ELEMENTARY CHEMISTRY 69 

under a pressure of 200 atmospheres. Finely divided iron is used 
as a catalytic agent. The ammonia formed is dissolved by water 
and the residual gases used over again (page 203). The method 
is used commercially in Germany as one of the ways for fixing 
nitrogen. Either ammonium sulfate for use as a fertilizer or 
nitric acid for use in making explosives is made from the 
ammonia. 

Laboratory preparation. On a small scale or for laboratory 
purposes, ammonia is prepared from ammonium salts, such as 
ammonium sulfate ((NH 4 ) 2 80 4 ) or ammonium chloride (NH 4 C1), 
by treating them with a base, like sodium hydroxide (NaOH) or 
calcium hydroxide (Ca(OH) 2 ). Ammonium hydroxide (NH 4 OH) is 
formed, which is unstable and decomposes to form ammonia and 
water. For a description of the method, see page 202. The 
reaction is a simple double decomposition in its first step: 

NaOH+NH 4 Cl-*NaCl+NH 4 OH 
NH 4 OH->NH 3 +H 2 

Properties. Consult page 203 for the physical properties of 
ammonia. Its odor is very striking and the gas can be unmis- 
takingly recognized by the sense of smell. It is very soluble in 
water. The experiment described on page 170 in connection with 
hydrogen chloride works equally well with ammonia. The density 
of the ammonia solution is less than that of water as there is 
expansion when the gas is dissolved. The great solubility of 
ammonia may be partly explained by the fact that some of it 
combines chemically with the water. 

The important chemical actions of ammonia are not many. 
Of greatest importance is the one with water just mentioned. 
Most of the common reactions which we think of in connection 
with ammonia are, in reality, the reactions of its compound with 
water, ammonium hydroxide. Other reactions of ammonia are: 
with metals, like magnesium, it will form nitrides, for example, 
(Mg 3 N 2 ); and it will burn with oxygen at high temperatures. 

Ammonium hydroxide. When water and ammonia combine, 
ammonium hydroxide (NH 4 OH) is formed. The solution is found 
to have basic properties, such as turning litmus blue and neutraliz- 
ing acids. This can be so only if OH~ ions are present in the 



70 ELEMENTARY CHEMISTRY 

solution. The NH 4 OH molecule is supposed to be formed and to 
dissociate into NH 4 + and OH~. Ammonium hydroxide has never 
been separated from its solutions. The radical (NH 4 ) forms the 
cation and otherwise has metallic properties; hence the name 
"ammonium" is given to it. A well-defined series of salts are 
formed with all the well-known acids, and these are called ammo- 
nium salts. Ammonium sulfate (NH 4 ) 2 S0 4 , ammonium chloride 
(NH 4 C1), and ammonium nitrate (NH 4 N0 3 ) are examples. 

Uses. Ammonia is used for making ammonium salts, which 
are used in fertilizers. The expansion of liquid ammonia is 
utilized in making artificial ice (page 110). 

Composition. The volume composition of ammonia is shown 
by decomposing a known volume of ammonia gas into hydrogen 
and nitrogen. It is found that 2 volumes of ammonia give 1 
volume of nitrogen and 3 volumes of hydrogen. 

Nitrogen acids. Of the nitrogen acids, nitric acid (HN0 3 ) is 
by far the most important. Another is nitrous acid (HNO2), 
which is unstable and weak; its sodium salt is used in dye manu- 
facture. These compounds do not occur in nature, but are made 
from other compounds. 

Preparation of nitric acid. Most of the nitric acid used is 
made from sodium nitrate (NaN0 3 ), commonly called Chili salt- 
peter. The method and the reaction are described on page 207. 
The salt is treated with concentrated sulfuric acid and heated 
gently; no reaction seems to take place, if the mixture is left cold. 
If dilute acid is used, mainly water distills off; but under proper 
conditions, nitric-acid vapor distills over and condenses. The 
explanation of this reaction will be taken up under the subject of 
equilibrium in the next lesson. 

Depending upon the amount of sulfuric acid used, the reac- 
tion is represented in two ways. Usually sodium acid sulfate is 
formed, as in the following equation: 

NaN0 3 +H 2 S0 4 ^NaHS0 4 +HN0 3 

With less sulfuric acid and more heat, normal sodium sulfate is 
formed, but the higher temperature decomposes some of the nitric 
acid, so the process is wasteful: 

2NaN0 3 +H 2 S0 4 -+Na 2 S0 4 -f2HN0 3 



ELEMENTARY CHEMISTRY 71 

The same reactions are made use of in preparing nitric acid com- 
mercially as in the laboratory. 

In order that the nitrogen of the atmosphere may be utilized 
to make nitric acid, other methods have been devised. One 
brings about the union of a part of the oxygen and the nitrogen 
of the air to form nitric oxide. This requires a high temperature, 
over 2000° C, which is produced by a magnified electric arc 
(Fig. 87, page 209). The oxide is passed into water, where a 
dilute solution of nitric acid is formed. The method is known as 
the Birkeland-Eyde process. The acid is neutralized with lime, 
forming calcium nitrate (Ca(N0 3 )2). This is known as air salt- 
peter and is used as a fertilizer. 

Nitric acid can also be formed by the oxidation of ammonia 
in the presence of a catalyzer. These two methods were used 
during the war by Germany to make nitric acid, as they could not 
import Chili saltpeter, because of the naval blockade. 

Properties of nitric acid. The physical properties of nitrir 
acid are set forth sufficiently on page 209. Its chemical conduct 
may be best considered under two general kinds of chemical 
behavior: (1) as an acid; (2) as an oxidizing agent. Its behavior 
on heating and its action with metals belong under one or the 
other of these heads. 

1. As an acid. Nitric acid is a strong acid, that is, in mod- 
erately dilute solutions it is highly ionized and conducts the 
electric current. It has a sour taste, turns litmus red, and 
neutralizes bases, forming salts called nitrates. With certain very 
strong metals and very dilute solutions of the acid, it can have 
its hydrogen replaced, thus: 

Mg+2HN0 3 ->Mg(N0 3 ) 2 -fH 2 

Ordinarily, however, any hydrogen that may be replaced is acted 
upon by other molecules of nitric acid, reduction products of 
nitric acid being formed (page 211). 

2. As an oxidizing agent. Nitric acid is a strong oxidizing 
agent because of its easy decomposition. In the presence of 
reducing agents, the oxygen which is available for oxidizing pur- 
poses is shown by the following equation: 

2HN0 3 ->H 2 0+2NO+3[0] 



72 ELEMENTARY CHEMISTRY 

The oxygen is enclosed in brackets to show that it is not actually 
set free. This oxygen will combine with the reducing agent, 
carbon, hydrogen, or a metal as follows: 

C+2[0]->C0 2 
H 2 +[0]-*H 2 
With most metals the oxide, which is formed, is basic in its 
properties and combines with more nitric acid to form the nitrate 
of the metal. The action of metals above hydrogen in the 
electrochemical series, such as zinc or iron, is best represented as 
shown on page 211. For metals after hydrogen, such as copper or 
silver, the equation near the middle of page 212 best represents 
what takes place. In the decomposition of nitric acid on heating, 
we may imagine that it first decomposes in the same way as 
when it acts as an oxidizing agent, but that in the absence of a 
reducing agent some of the oxygen combines with the nitric oxide 
(NO) to form dioxide (N0 2 ) as follows: 

4HN0 3 ->2H 2 0+4NO+6[0] 

4NO+6[0]->4N0 2 +0 2 
Adding we have ■■ ^ 

4HN0 3 -*2H 2 0+4N0 2 +0 2 

which represents the final results of the action. Nitric acid will 
oxidize certain compounds by removing hydrogen from them. 
An example is the action between hydrochloric and nitric acids. 
These equations show the result of the action: 
2HN0 3 ^H 2 0+2NO+3[0] 
6HCl+3[0]->3H 2 0+3Cl 2 
This mixture of acids is called aqua regia (page 213) because of its 
power to dissolve gold, the king of metals. It is the nascent 
chlorine that is the active agent. 

Oxides of nitrogen. There are five oxides of nitrogen. Their 
names, formulas, and physical states are given in the table on 
page 214. Three can easily be prepared. Two, nitrogen trioxide 
and nitrogen pentoxide, are hard to make and are only important 
as they are related to nitrous and nitric acids. When added to 
water they combine according to the following equations: 

N 2 3 +H 2 0->2HN0 2 

N 2 5 +H 2 0->2HN0 3 
For this reason they are called acid anhydrides (page 217). 



ELEMENTARY CHEMISTRY 73 

Anhydrides. The term "anhydride" may be given a broader 
application and made to include basic oxides as well as acid 
oxides (see "Introduction," Lesson X). An anhydride may be 
defined as an oxide which can be obtained from an acid or a base by 
the elimination of water. Examples of acid anhydrides and their 
acids are: S0 2 -*H 2 S0 3 , C1 2 0->2HC10, S0 3 ->H 2 S0 4j and those 
given in the previous paragraph. Examples of basic anhydrides 
and their bases are: K 2 0->2KOH, Na 2 0->2NaOH, CaO->Ca(OH) 2 , 
Fe 2 3 ->2Fe(OH) 3 . 

Nitrous oxides. Nitrous oxide is frequently called laughing 
gas and is used by dentists in extracting teeth, as it is a mild 
anesthetic. It is prepared by heating ammonium nitrate: 

NH4N0 3 ->2H 2 0+N 2 

It will support combustion easily as it is sufficiently unstable to 
give up its oxygen. 

Nitric oxide. Nitric oxide is made by the action of nitric 
acid on a metal; copper is usually used. The reactions are those 
given under nitric acid on page 212. It is a colorless gas which 
unites readily with oxygen to form nitrogen dioxide (N0 2 ) (page 
215). It is more stable than nitrous oxide and will not support 
the combustion of wood or sulfur, but it will support the com- 
bustion of phosphorus. 

Nitrogen dioxide. Nitrogen dioxide can be made by decom- 
posing nitrates (page 216). It is a brown gas with suffocating 
odor and is a poison. It gives up part of its oxygen and acts as 
an oxidizing agent. 

EXERCISES 

1. How was ammonia originally prepared? 

2. Give the usual commercial method for making ammonia. 

3. State the essentials of the Haber process for making ammonia. 

4. Give the laboratory method for making ammonia. Give equations. 

5. What property characterizes compounds of nitrogen? 

6. Give the physical properties of ammonia. 

7. State the chemical reactions of ammonia. 

8. What is the ammonium radical? Give its properties. 

9. State the uses of ammonia. 

10. What is the volume composition of ammonia? 

11. Give the usual method for preparing nitric acid. Equations. 

12. Give two other methods for making it. Describe the Birkeland- 
Eyde method. 



74 ELEMENTARY CHEMISTRY 

13. What are the physical properties of nitric acid? 

14. Give a full statement of the chemical properties of nitric acid. 

15. Give name, formula, and a distinguishing property for the oxides of 
nitrogen. 

16. How is laughing gas made? 

17. Give its properties and uses. 

18. How are nitric oxide and nitrogen dioxide made? 

19. Define anhydride. Give three examples each of acid and basic 
anhydrides. 

20. Write equations for the action of nitric acid on mercury. 

21. In what different ways do metals act on nitric acid? 

22. What is aqua regia? Write the equation for its action. 

23. Answer questions 1, 2, 3, 4, 5, 6, 7, and 8, page 217. 

24. Solve problems 14, 15, 16, 17, 18, and 19, page 218. 

25. Give chemical name, formula, and use for air saltpeter, Chili salt- 
peter, and laughing gas. 

LESSON XIII 

EQUILIBRIUM; PERIODIC SYSTEM 
EQUILIBRIUM 
Assignment: Chapter XIX, McPherson and Henderson 

Speed of a reaction. Some reactions take place with a much 
greater velocity than others. Most inorganic reactions proceed 
very rapidly, while organic reactions are generally slow. How- 
ever, iron rusts slowly. At ordinary temperature copper and 
oxygen combine slowly, and the action of pure zinc on sulfuric 
acid is very slow. By the speed of a reaction is meant the 
amount which undergoes change in a given time. 

Besides the nature of the substance, three factors influence 
the velocity of a reaction. 

1 . Temperature. Raising the temperature promotes the speed 
of a reaction. Coal burns faster when the fire is hot. A metal 
acts faster with an acid if we heat it. Cold storage retards the 
chemical change of decay. "Pressure cookers/ ' which give a tem- 
perature higher than usual for boiling water, shorten the time of 
cooking. 

2, Concentration. It was seen that substances burn faster 
in pure oxygen than in air. Zinc will dissolve faster in concen- 
trated hydrochloric acid than in dilute. Instead of increasing the 
concentration, the same result may be obtained by increasing the 



ELEMENTARY CHEMISTRY 75 

surface of the reacting substances. Coal dust and oxygen com- 
bine so rapidly as to cause explosion at times. Kindling burns 
faster than the same wood in logs. An excess of one substance 
causes no more action, but it does cause it to go faster. Forcing 
air over coal will make it burn faster. The use of an excess of 
one substance is to produce the effect of mass action. 

3. Catalysis. Sometimes certain substances can be added to 
a reaction and cause it to go faster or slower without seeming to 
take any part in it. At least, the added substance comes out of 
the reaction the same in kind and quantity as when it went in. 
These substances are called catalyzers, and the effect produced is 
known as catalysis. We have had examples in the preparation of 
oxygen (page 26) and of ammonia (page 203). Water is some- 
times a catalyst; dry carbon monoxide (CO) will not burn. A 
class of substances, called enzymes, act by catalysis. Enzymes 
are formed by bacteria and promote such reactions as the split- 
ting of cane sugar into dextrose and fructose, the fermenting of 
dextrose to alcohol, the souring of milk, and the decay of organic 
matter generally. Enzymes play an important part in the pro- 
cesses of digestion. 

Reversible reactions. Several examples of reactions which 
will go in either direction according to conditions have already 
been met with. These are called reversible 'reactions. This prop- 
erty of the reaction may be represented by the double arrow in 
the equation. 

30 2 ^20 3 (page 114) 

2NH 3 ^tN 2 +3H 2 (page 203) 

2Hg0^2Hg+0 2 (page 25) 

Burning limestone gives lime and carbon dioxide, but lime will 
absorb carbon dioxide at ordinary temperatures: 

CaC0 3 ^CaO+C0 2 

Nearly all reactions are reversible, though not always in as simple 
a way as in the examples above. 

Equilibrium. The reaction of steam and heated iron is a 
reversible reaction. The products, iron oxide and hydrogen, will 
form steam and iron: 

3Fe+4H 2 C£±Fe 3 4 +4H 2 (page 40) 



76 ELEMENTARY CHEMISTRY 

If steam is passed over the iron and the hydrogen allowed to 
escape, the reaction goes from left to right. However, if we place 
iron oxide in the tube and pass hydrogen, the steam formed will 
escape and the reaction will go from right to left. 

Now imagine the molecular equivalent amounts of iron and 
steam sealed up in the tube and the tube heated. At first the 
reaction will be mostly from left to right. As the iron and 
steam are used up, this reaction slows down. Meanwhile, the 
reaction between iron oxide and hydrogen has been gaining 
velocity, as these substances were formed in greater quantities. 
This will continue until the speed of the reverse reaction equals 
the speed of the original reaction. This point is called chemical 
equilibrium. This point will be reached when the molecular 
quantities used up in one direction are just offset by the molecular 
quantities of the same substances formed by the reaction in the 
other direction. 

Chemical equilibrium is not to be considered as a state of 
rest, but, rather, as the balance between two opposite forces 
which exactly equal each other. If you are walking on a moving 
sidewalk in the opposite direction to which it is moving, but with 
the same speed, you appear at rest with reference to a lamp post 
by the side of the walk. It is possible to have all the substances 
involved in reversible reactions present together in equilibrium 
with each other. The average percentage of each material will 
then remain unchanged. In the formation of ammonia by the 
electric spark, equilibrium is reached when 7 per cent of ammonia 
is present. 

Increasing the mass of a substance on the left or removing 
one of the products on the right will result in a change in the 
equilibrium. In the example above, if we add more steam, or 
remove the hydrogen, or do both, the reaction will go more 
nearly to completion. In general, a reversible reaction may be 
made to go to completion if one of the products of the reaction is 
removed as fast as it is formed. 

Equilibrium in solution. Many of our most familiar and 
important reactions take place in solution. Here, we are con- 
cerned primarily with the equilibrium of ions with the molecules 
from which they are formed. Mixing the solutions of two com- 



ELEMENTARY CHEMISTRY 77 

pletely dissociated electrolytes does not . change the ionic condi- 
tion; we have all four kinds of ions present. Most solutions do 
not completely dissociate their electrolytes. In such cases we 
have all four kinds of molecules present, as well as all four kinds 
of ions. Take sodium chloride and potassium nitrate: 

NaCl+KNQ^KCl+NaNO, 

The original solutions contain the molecules NaCl and KN0 3 and 
the ions Na + , Cl~, K + and NO3 - . When mixed, these ions form 
the molecules KC1 and NaN0 3 as readily as the original NaCl 
and KN0 3 . 

Completion of reactions in solutions. Such double decompo- 
sition reactions in solutions will go to completion only when one 
of the products is removed from the reaction. This may happen 
only if one of the following conditions is established. 

1. A volatile gas may form. This includes cases where the 
product of the action is itself an insoluble gas or a very unstable 
substance like ammonium hydroxide which splits up to form a 
gas. Complete reaction results if conditions are changed so as 
to make a substance a gas which is not usually a gas or to pre- 
vent solution if the product is naturally soluble. The prepara- 
tion of nitric acid is explained on this basis. The mixture was 
heated to make nitric acid a gas and concentrated sulfuric acid 
was used to prevent the solution of the nitric acid in water: 

NaN0 3 +H 2 S0 4 ->NaHS04+HN0 3 

Other examples are: 

2NaCl+H 2 S0 4 -+Na 2 S0 4 +2HCl 

NaOH+NH 4 Cl->NaCl+NH 3 +H 2 

CaC0 3 +2HCl->CaCl 2 +C0 2 +H 2 

2. An insoluble solid may form. If an insoluble solid forms, 
the ions that form the insoluble molecule drop out of the solution 
as a precipitate and can take no part in the reverse reaction: 

AgNQ 3 +HCl->AgCl-fHNQ 3 
Pb(N0 3 ) 2 +H 2 S0 4 -^PbS04+2HN0 3 

The compounds underscored are the ones that precipitate. The 
others remain in solution partly ionized. 



78 ELEMENTARY CHEMISTRY 

3. An undissociated molecule may form. The prominent 
example in this case is the reaction of an acid and a base. All 
acids and bases react to form an undissociated molecule of water 
(water is very slightly ionized). For examples consult Lesson X. 

Hydrolysis. An example of a reversible reaction is the reac- 
tion of water with some salts. With a salt of a strong acid and 
base the reaction, NaCl+HOH— >NaOH+HCl, is not noticeable 
because both sodium hydroxide and hydrochloric acid are strong 
and the tendency is altogether in the opposite direction. If either 
the acid or base formed is weak the result is different. In the 
case of 

Na 2 C03+2HOH->2NaOH+H 2 C0 3 
the carbonic acid is weak. Its molecules do not form sufficient 
hydrogen ions to neutralize the hydroxyl ions of the base. There- 
fore, we have an excess of hydroxyl ions, and such a solution will 
turn litmus paper blue. 

On the other hand, if the base formed is weak, the opposite 
result will be obtained. A solution of ferric chloride will turn 
litmus red: 

FeCl 3 +3HOH-*3HCl+Fe(OH) 3 

The iron hydroxide is very weak, while the acid is strong. The 
reactions are called hydrolysis. Hydrolosis is the action of water 
upon a salt to form an acid and a base, one of which must be weak. 
If both acid and base are weak, the reaction may be complete 
hydrolysis. This is the case with aluminium carbonate and water: 

Al 2 (C0 3 )3+6HOH-^2Al(OH)3+3C0 2 +3H 2 
This occurs in the use of alum baking powders, the alum first 
forming the aluminium carbonate. 

PERIODIC SYSTEM 

Assignment: Chapter XXI, McPherson and Henderson 

Classification of elements. It was early recognized that 
some relation existed between the atomic weights of elements and 
their properties. Dobereiner pointed out that several groups of 
three elements exist which were closely similar in properties and 
whose atomic weights were in arithmetical proportion. For exam- 
ple, chlorine, bromine, and iodine have quite similar properties. 



ELEMENTARY CHEMISTRY 79 

The atomic weight of bromine is an approximate mean between 
those of chlorine and iodine: 

35.46+126.92-5-2 = 81.19. Bromine is 79.92 

Again, a group calcium, strontium, and barium give similar 
results : 

40.07+137.37^2 = 88.72. Strontium is 87.63 

These groups were called Dobereiner's triads. 

The triads were incomplete, including only a part of the 
elements. In 1866 Newlands suggested his octaves. He arranged 
the elements in the order of increasing atomic weight, starting 
with hydrogen, and pointed out that every eighth element showed 
a similarity of properties. In 1869 Mendeleeff and Lothar Meyer 
brought out their periodic tables which are essentially the same 
and are a more complete elaboration of the idea of Newlands. 

Periodic System. The table most frequently used and the 
one shown on page 258 is that of Mendeleeff brought up to date. 
Hydrogen is omitted as not having properties which properly 
place it with any of the groups. There are eight elements before 
properties begin to repeat. The elements found in any vertical 
column have marked similarity of properties. Lithium, sodium, 
potassium, etc., in Group I are much alike. So are fluorine, chlo- 
rine, bromium, and iodine in Group VII. Helium, neon, argon, 
etc., are placed in Group O because they were not known when 
Mendeleeff numbered the other groups. 

It will be noticed that some elements in the same group are 
more similar to each other than they are to the others. Each 
group is therefore divided into two families. One family is 
placed to the left and the other to the right of the group. In 
Group I sodium should be placed with the left-hand group, as 
its properties are much more like those of lithium, potassium, 
rubidium, and caesium than they are like those of copper, silver, 
and gold. In Group VII fluorine had better be placed with the 
chlorine family. 

It is seen that the properties of the elements do not vary 
continuously but repeat at regular intervals, or periods. The 
periodic law is stated thus: The properties of the elements are 
periodic functions of the atomic weights. 



80 ELEMENTARY CHEMISTRY 

Uses of system. This arrangement of the elements has 
served to predict new elements. Mendeleeff predicted the exist- 
ence of scandium, germanium, and gallium. The vacancies in the 
table may be filled by new discoveries. It has served to make 
possible a decision between two possible atomic weights for an 
element. But, for us, its most valuable use is that it simplifies 
study. A study of the properties of one element of a family will 
make easier the study of the others. 

The physical as well as the chemical properties of the mem- 
bers of a family will be much the same. Members of the same 
family have the same valences. With hydrogen, the valence of 
the group increases from zero to four and then decreases. With 
oxygen there is an increase from zero to eight, though eight is 
not the common valence of the elements of Group VIII. In the 
following lessons we will study the elements in their families and 
point out in detail the application of the periodic law to the 
study of the properties of the elements. 

EXERCISES 

1. What is meant by the speed of a reaction? 

2. Name the factors that influence speed. Give examples of each. 

3. Define catalysis; enzyme. Give examples. 

4. What is a reversible reaction? Give several examples. 

5. Explain the effect of mass on the velocity of a reaction; on chemical 
equilibrium. 

6. Define chemical equilibrium. Give an example. 

7. State the general condition necessary for a reversible reaction to go 
to completion. 

8. State the condition of equilibrium in a mixed solution of two elec- 
trolytes, all substances being soluble. 

9. Name three conditions under which a double decomposition reaction 
will go to completion in solution, using equations. Give two examples of each. 

10. Explain the preparation of nitric acid. 

11. Define hydrolysis. Give two examples. 

12. Answer questions 3, 4, 5, 6, and 9, pages 226 and 227. 

13. State the periodic law. 

14. What were Dobereiner's triads? Newlands' octaves? 

15. Briefly state the arrangement of the elements in the periodic table. 

16. What are the uses of the periodic system? 

17. Answer questions 1, 2, 5, and 6, page 263. 



ELEMENTARY CHEMISTRY 81 

LESSON XIV 

SULFUR AND ITS COMPOUNDS 
Assignment: Chapter XX, McPherson and Henderson 

Occurrence. Sulfur occurs both free and in the combined 
state. Free sulfur is found in two kinds of deposits in nature. 1 
Sulfur found in certain volcanic regions in Sicily and Japan until 
recently formed the world's principal supply. These deposits are 
formed from the gases emitted during an eruption; hydrogen sul- 
fide (H 2 S) is present and is partially burned, the sulfur being 
deposited. 

In the United States sulfur is found in Louisiana far under- 
ground. It is obtained by forcing superheated water into wells 
drilled into the deposits. This melts the sulfur, which is forced 
out of another pipe by compressed air. The liquid sulfur then 
solidifies. This source of sulfur now supplies the needs of this 
country and much is exported. 

The sulfur deposits in Louisiana have an interesting origin. 
It is supposed that beds of calcium sulfate (gypsum) first existed. 
This was reduced to calcium sulfide by decaying organic matter. 
At the same time acids were formed which changed the calcium 
sulfide into hydrogen sulfide. Certain species of algae (a form of 
plant life) consumed the hydrogen sulfide, and, in the course of 
their life processes oxidized it to sulfuric acid. This in turn con- 
verted the calcium salts back to calcium sulfate. The cycle of 
sulfur was thus completed. But, the algae consumed more hydro- 
gen sulfide than could be completely oxidized to sulfuric acid, 
and this surplus was changed to water and sulfur. The sulfur 
was stored up in the plant cells, much as other plants store up 
starch or animals store up fat. In the course of time these plants 
died, and their bodies decayed, leaving the deposit of sulfur. 

Allotropic forms of sulfur. Four allotropic forms of sulfur 
are common; two of these are crystalline and two amorphous. 

/. Rhombic sulfur. Rhombic sulfur is the ordinary form; 
brimstone is chiefly rhombic sulfur. This form is obtained by 
crystallizing sulfur from carbon disulfide solution. It appears as 
yellow plates usually with eight sides. 



82 ELEMENTARY CHEMISTRY 

2. Monoclinic sulfur. When sulfur is melted and allowed to 
cool, monoclinic sulfur is obtained. The crystals are needle 
shaped. Rhombic sulfur changes to monoclinic sulfur ab@ve 95. &° 
C, which is known as the transition temperature; below thk 
temperature the change is reversed. 

3. Plastic sulfur. Plastic sulfur is amorphous and elastic, 
like rubber. It quickly passes into the rhombic form. It can be 
obtained by suddenly cooling sulfur in water. 

4. "Flowers of sulfur." When sulfur first passes into the 
cold, some of it condenses to a fine crystalline powder, known as 
"flowers of sulfur." 

These four forms are considered to be allotropic for the same 
reasons as those given for so considering the forms of carbon. 

Properties and uses. The physical properties vary with tne 
different forms. Sulfur burns to form a gas, sulfur dioxide (S0 2 ). 
It will combine with most metals, forming sulfides. It is gener- 
ally similar to oxygen in its chemical activity; it is not so active 
as oxygen is, but combines with about the same number of 
elements. 

Sulfur is used in the manufacture of gunpowder, vulcanized 
rubber, lime sulfur spray, carbon disulfide, sulfur dioxide, and 
sulfuric acid. Lime sulfur spray is used as an insecticide for 
spraying fruit trees. 

Hydrogen sulfide. This compound of hydrogen and sulfur is 
a disagreeable-smelling gas, the odor being suggestive of badly 
decayed eggs. It is formed in the decay of albuminous matter. 
It is found in volcanic gases and in certain waters, known as 
"sulfur waters." It can be made by the direct union of the two 
elements, but it is usually made from a metallic sulfide and an 
acid. Iron sulfide is used: 

FeS+2HCl-»FeCl 2 +H 2 S 

This reaction goes to completion because hydrogen sulfide is 
much less soluble in water than the hydrogen chloride. The water 
soon becomes saturated with hydrogen sulfide, and the excess of 
the gas escapes (page 233). 

Properties of hydrogen sulfide. The physical properties of 
this substance can be studied on page 233. A study of its chem- 



ELEMENTARY CHEMISTRY 83 

ical conduct shows that its chemical properties can be best con- 
sidered as of two kinds: 

1. Acid properties. In water solution it is a weak acid. It 
is sometimes called hydrosulfuric acid. It has the properties of 
an acid, neutralizing bases, forming salts, and turning litmus red. 
Its salts are called sulfides. 

2. Reducing action. Hydrogen sulfide is an unstable sub- 
stance. When heated to 500° C, the speed of decomposition is 
very noticeable. Slow decomposition into hydrogen and sulfur 
takes place in solution. Because of this unstability it will react 
with metals below hydrogen in the electrochemical series, which, 
as an acid, it would not do. Silver is an example. Silver spoons 
and other tableware are tarnished black by the hydrogen sulfide 
in the air, especially when the cheaper grades of soft coal (which 
contain much sulfur) are burned. The ease with which hydrogen 
sulfide gives up hydrogen makes it a good reducing agent; the 
compound "will combine readily with oxygen. With sufficient 
oxygen and high temperature we get 

2H 2 S+30 2 ->2H 2 0+2S0 2 

With less oxygen water and sulfur are formed: 

2H 2 S+0 2 ->2H 2 0+2S 

It will reduce certain oxidizing agents by taking the oxygen from 
them. With dilute nitric acid the action is shown thus: 

2HN0 3 ->H 2 0+2NO+3[0] 

3H 2 S+3[0]->3H 2 0+3S 

Summing up the two steps we have 

2HN0 3 +3H 2 S->4H 2 0+2NO-f3S 

The sulfides. One of the principal uses of hydrogen sulfide 
is as a reagent to form sulfides of the metals. Most of the 
metallic sulfides are insoluble in water. The insoluble ones can 
be grouped into two classes: those insoluble in dilute acids, and 
those insoluble in alkaline solution. In this way the metals are 
divided into groups for the purposes of analysis. The sulfides 
which are insoluble can be prepared by passing hydrogen sulfide 



84 ELEMENTARY CHEMISTRY 

through solutions of the salts of the proper metals. The soluble 
ones are made by neutralizing the proper base with hydrogen 
sulfide. 

Sulfur dioxide. There are two oxides of sulfur, sulfur dioxide 
(SO2) and sulfur trioxide (S0 3 ). They are the anhydrides of 
sulfurous acid (H 2 S0 3 ) and sulfuric acid (H 2 S0 4 ). Sulfur dioxide 
is the product of burning sulfur in oxygen; it has an irritating 
odor; under proper conditions it combines with oxygen to form 
sulfur trioxide; it may act as a reducing agent: 

2H 2 S+S0 2 ->2H 2 0+3S 

Sulfur dioxide combines with water to form sulfurous acid, and 
many of the reactions assigned to it are really the reactions of 
sulfurous acid (page 238). 

Preparation of sulfur dioxide. It can be made in three ways: 

1. Burning sulfur or a sulfide. The reaction is as follows: 

2ZnS+30 2 ->2ZnO+2S0 2 

2. Reducing sulfuric acid. The following equations show 
what happens: 

H 2 S0 4 ->H 2 S03+[0] 

Cu+[0]->CuO 

CuO+H 2 S0 4 ->CuS0 4 +H 2 

H 2 S0 3 ->H 2 0+S0 2 

These steps give on addition 

Cu+2H 2 S0 4 ->2H 2 0+S0 2 +CuS0 4 

3. Action of an acid on a sulfite. Sulfurous acid is first 
formed. It is unstable and breaks up to give water and sulfur 
dioxide : 

Na 2 S0 3 +2HCl->2NaCl+H 2 S0 3 
H 2 S0 3 ->H 2 0+S0 2 

This reaction would be reversible and establish equilibrium, except 
that the sulfurous acid decomposes as fast as it forms, giving 
sulfur dioxide, which, as a gas, escapes. For this reason the action 
is practically complete. 

Sulfurous acid. Sulfurous acid is so unstable that it has 
never been isolated from the solution in which it is obtained 
when sulfur dioxide is passed into water. Some sulfurous acid 



ELEMENTARY CHEMISTRY 85 

molecules are formed and they ionize to form hydrogen ions." This 
is shown by the acid properties of the solution. Heating, how- 
ever, drives off the sulfur dioxide, as if it were held in physical 
solution. Sulfurous acid has the properties^ of a weak acid and 
forms salts known as sulfites, "f] 

It is a reducing agent, having the power to take up another 
atom of oxygen to form sulfuric acid. It will reduce nitric acid: 

2HN0 3 ->H 2 0+2NO+3[0] 
3H 2 S0 3 +3[0]-»3H 2 S0 4 
or 

2HN03+3H 2 S0 3 ->H 2 0+2NO+3H 2 S0 4 

It has bleaching and antiseptic properties and is used to bleach 
straw and flour. It is sometimes used as a preservative in cer- 
tain foods. Calcium acid sulfite is used in paper manufacture 
(page 240). 

Sulfur trioxide. Only traces of sulfur trioxide (S0 3 ) are 
obtained when sulfur burns or when sulfur dioxide and oxygen 
are heated together. In order to get sulfur trioxide in quantity, 
a catalytic agent must be used. When sulfur dioxide and oxygen 
are passed over finely divided platinum at 400°, 98 per cent of 
the sulphur dioxide combines with oxygen (page 241). The prin- 
cipal use of sulfur trioxide is to make sulfuric acid by the con- 
tact process. It combines readily with water, forming the acid. 
For other properties see page 242. 

Sulfuric acid. Sulfuric acid (H 2 S0 4 ) is one of the most 
important substances used by chemical industries. It is used in 
the manufacture of explosives, fertilizers, and dyes, in the refining 
of petroleum, and in the making of other acids. Sulfuric acid 
cannot be made in the manner that nitric and hydrochloric acids 
are made, because of its high boiling point. 

Manufacture of sulfuric acid. Two methods for making sul- 
furic acid are in common use, the contact process and the lead- 
chamber process. The contact process is described on page 243. 
It is gradually being used in this country for the purpose of 
making concentrated sulfuric acid. The chamber process is the 
older method and is still used. A description of the process is 
found on pages 244, 245, and 246. 



86 ELEMENTARY CHEMISTRY 

In studying, the student should recognize two distinct parts 
to the process: (1) the making of the acid, including the gener- 
ation of sulfur dioxide and the introduction into the chambers of 
oxides of nitrogen, steam, and air and the reactions resulting; (2) 
that part of the process which is designed to save the oxides of nitro- 
gen for repeated use. The Glover and Gay-Lussac towers, com- 
monly called the hot and the cold towers, are used for the latter 
purpose, the oxides of nitrogen being dissolved in the cold tower 
and liberated in the hot tower, from which they pass with the 
sulfur dioxide into the lead chambers for use again. 

This process gives an acid of about 65 per cent pure hydrogen 
sulfate. It is the cheaper method for making dilute sulfuric acid, 
while the contact process can compete with it in making the con- 
centrated acid. Chamber acid can be concentrated by evapora- 
tion in platinum dishes. 

Properties of sulfuric acid. Sulfuric acid is a heavy oil, 
sometimes called oil of vitriol. It has a high boiling point, 338° 
C. for the ordinary concentrated acid, which contains 2 per cent 
water (page 246). Its chemical behavior can best be discussed 
under four kinds of actions: 

1, Acid properties. In water solution it is highly disso- 
ciated into ions and has the properties of a strong acid. It 
forms the ions H + and HSO~ 4 in concentrated solution and H + , 
H + , S0 4 ~ ~ in dilute solution. It reacts with bases to form normal 
salts (Na 2 S0 4 ) and acid salts (NaHS0 4 ). With metals above 
hydrogen in the electrochemical series, the dilute acid will act to 
form hydrogen and a sulfate of the metal. 

2. Oxidizing action. When concentrated sulfuric acid is 
heated, it acts as a good oxidizing agent, like, but not so good as, 
nitric acid. One molecule of the acid will give up one atom of 
oxygen to combine with the reducing agent: 

H 2 S0 4 ->H 2 0+S0 2 +[0] 

The reducing agent may be carbon, hydrogen, a metal, or some 
compound. These equations show its action with carbon: 

2H 2 S0 4 -»2H 2 0+2S0 2 +2[0] 

C+2[0]-»C0 2 



ELEMENTARY CHEMISTRY 87 

Adding, the final result is 

2H 2 S0 4 +C^2H 2 0+2S0 2 -r-C0 2 
Equations with copper are given on page 247. With aluminium 
and hot concentrated sulfuric acid, the following equations rep- 
resent the action: 

3H 2 S0 4 ->3H 2 0+3S0 2 +3[0] 
2Al+3[0]->Al 2 3 
A1 2 3 +3H 2 S0 4 ^A1 2 (S0 4 ) 3 +3H 2 

Combining these steps we have 

2A1+6H 2 S0 4 ^6H 2 0+3S0 2 +A1 2 (S0 4 ) 4 

3. Action with salts. This action depends upon the fact 
that sulfuric acid has a higher boiling point than the other com- 
mon acids. When sulfuric acid is added to salts at a tempera- 
ture above the boiling points of their acids, the acids escape as 
gases (page 225). 

4> Action on water. Concentrated sulfuric acid is a good 
dehydrating agent because of its ability to combine readily with 
water. It is used as a drying agent because it will absorb moist- 
ure from gases that do not act with it. 

Basicity of acids. Acids which form but one hydrogen ion 
will neutralize but one unit of a base and are called monobasic 
acids. Nitric acid (HN0 3 ) is a monobasic acid. Acids, like 
sulfuric acid (H 2 S0 4 ), which form two hydrogen ions will neu- 
tralize two units of a base. They are called dibasic acids. We 
also have tribasic (H 3 P0 4 ) and tetrabasic (H 4 SiQ 4 ) acids. 

Carbon disulfide. Carbon disulfide is a heavy, disagreeable- 
smelling, colorless, oily liquid. It is said to have a pleasant 
odor when pure. It is used as a solvent for sulfur, rubber, gums, 
resins, and waxes. It is also used as an insecticide, its vapor 
being poisonous. It is made by heating carbon and sulfur together 
in an electric furnace (page 250). 

Hydrates. The compounds which are formed when sulfuric 
acid, for example, combines with water are known as hydrates. 
Many salts have the property of combining with water in a 
similar way. Copper sulfate combines with five molecules of 
water to form the blue vitriol crystals (CuS0 4 .5H 2 0). Similarly we 
have such hydrates as Na 2 CO 3 .10H 2 O, BaCl 2 .2H 2 0, CaCl 2 .6H 2 0, 



88 ELEMENTARY CHEMISTRY 

and FeS0 4 .7H 2 0. Such substances are true chemical compounds 
since they have a fixed percentage composition. 

However, they are not very stable, but lose their water when 
heated slightly above the boiling point of water. Hydrates are 
compounds in which water molecules are held in a loose state of 
chemical combination with the main molecule. In dry air the water 
escapes at a lower temperature from some, like Na 2 SO4.10H 2 O. 
Such salts are said to effloresce. Efflorescence is the tendency of 
hydrates to lose water at ordinary temperatures. On the other hand, 
some salts, like calcium chloride (CaCl 2 ), tend to absorb water 
from the air to form hydrates. They may even dissolve in the 
water and become liquid. They are said to deliquesce. Deli- 
quescence is the tendency of salts to absorb water and become hydrates. 

EXERCISES 

1. Discuss the conditions under which sulfur occurs in nature. Explain 
its formation. 

2. Give the formation and properties of the allotropic forms of sulfur. 

3. State the chemical properties of sulfur. 

4. Name the uses of sulfur. 

5. Give the occurrence and preparation of hydrogen sulfide. Give the 
equation. 

6. Explain why the reaction for preparing hydrogen sulfide goes to 
completion. 

7. State the physical properties of hydrogen sulfide. 

8. What is hydrosulfuric acid? State its chemical properties. 

9. Write equations for three ways for preparing sulfur dioxide. 

10. Give the properties of sulfur dioxide. 

11. Discuss the properties of sulfurous acid. 

12. Give the preparation of sulfur trioxide. 

13. Describe the contact process for making sulfuric acid. 

14. Describe the lead-chamber process for making sulfuric acid. 

15. State the physical properties of sulfuric acid. 

16. Discuss fully the four kinds of chemical properties of sulfuric acid, 
using equations to illustrate, when possible. 

17. Define dibasic acid, tribasic acid, hydrate, deliquescence, efflores- 
cence. Give an example of each. 

18. How is carbon disulfide made? Give its properties and uses. 

19. Complete the following equations: 

Cu+hot concentrated H 2 S0 4 ->? ZnS+HCl— ►? 

H 2 S+dilute HN0 3 -+? Cu+dilute H 2 S0 4 -»? 

H 2 S+excess 2 ->? H 2 S+S0 2 ->? 

AsCl 3 +H 2 S->? H 2 S0 3 +HN0 3 -*? 

20. Answer questions 1, 2, 3, 4, 5, 6, 7, 8, and 9, page 252. 

21. Solve problems 11, 12, 13, 14, and 17, page 253. 



ELEMENTARY CHEMISTRY 89 

LESSON XV 

THE HALOGEN FAMILY 

Assignment : Chapter XXII and Review of Chapter XIV, McPherson 
and Henderson 

The family. The four elements, fluorine, chlorine, bromine, 
and iodine are placed in the second family of Group VII of the 
periodic table. They offer a good opportunity to observe the 
way in which members of the same periodic family both resemble 
and differ from each other. It will be noted that they differ 
only in the degree of their action, and not in the kind. 

Furthermore, their properties, physical and chemical, are in 
direct relation to their atomic weights. They are called halogens 
because they are foimd mostly in combination as salts in sea 
water. The word "halogen" means salt former. They do not 
occur in the free state because they are so active chemically 
that if ever liberated as free elements, they always find some 
other element to combine with. 

Physical properties. The table at the top of page 264 shows 
the principal physical properties of these elements. Note that 
there is an increase in melting point, in boiling point, and in 
darkness of color as the atomic weight becomes larger. Fluorine 
is a pale-yellow gas, chlorine a greenish-yellow gas, bromine a 
red liquid easily volatile, and iodine a black solid easily con- 
verted to a purple vapor. 

Chemical Conduct. These elements are found in the last 
regular group of the periodic table, far removed from the strong 
metals. If the electrochemical series should be extended to 
include the non-metals, the halogens, with oxygen, would be 
found at the very end. Therefore, they are strongly negative 
elements in their chemical behavior, in fact, the most active 
of all the elements in a negative way. 

As negative elements the halogens have properties similar to 
those of oxygen. For this reason they are very good oxidizing 
agents. The term "oxidation" must be thought of in a broader 
sense than in connection with oxygen alone; an oxidizing element 
is one ivith properties similar to oxygen. In this sense the halogen 



90 ELEMENTARY CHEMISTRY 

elements are very active oxidizing agents. Sometimes they 
remove hydrogen from water, freeing the oxygen, in which reac- 
tion they are indirect oxidizing agents (see "Bleaching action," 
page 164); other times they oxidize directly, as: 

2FeCl 2 +Cl 2 -*2FeCl 3 

It needs to be emphatically stated that all four of the halogen 
elements are good oxidizing agents. Some, however, are better 
than others. Their power as oxidizing agents, as well as their 
general chemical activity, varies inversely with their increasing 
atomic weights. Fluorine is most active, chlorine next, bromine 
next, and iodine least active as an oxidizing agent. Their tend- 
ency to form compounds with hydrogen and the metals varies 
in the same order. So does their bleaching action. With the 
other negative, or non-metallic, elements the halogens form less 
stable compounds. With oxygen only chlorine and iodine form 
compounds, and these are very unstable. Sometimes one halogen 
forms compounds with another halogen, but such compounds 
are very unstable. 

Compounds with hydrogen. All the halogen elements form 
compounds with hydrogen in which they have a valence of one. 
The hydrogen compounds, in water solution, are all acids and 
form a series of salts in which the halogens combine with a valence 
of one with the metals, forming fluorides, chlorides, bromides, and 
iodides. The halogen hydrides vary in stability as the activity 
of the halogen itself varies; that is, the more active element will 
form the more stable hydride. In the order of their stability they 
are: hydrogen fluoride, hydrogen chloride, hydrogen bromide, and 
hydrogen iodide. 

Since they are sufficiently unstable to give up their hydrogen 
under suitable conditions, these hydrogen compounds of the halo- 
gens are all reducing agents, some better than others. The best 
reducing agent is hydrogen iodide, the one with the largest molecular 
weight and the least stability, iodine being the least active halo- 
gen. Then comes hydrogen bromide, which is a better reducing 
agent than hydrogen chloride. It must be emphasized here that 
it is the hydrogen compounds of the halogens which act as reducing 
agents, the halogen elements being oxidizing agents. 



ELEMENTARY CHEMISTRY 91 

Hydriodic acid and hydrobromic acid will reduce both nitric 
and sulfuric acids, hydriodic acid carrying the reduction of sul- 
furic acid, at least in part, to sulfur and hydrogen sulfide, while 
hydrobromic acid reduces it only to sulfur dioxide. Hydrochloric 
acid will reduce nitric acid (see the discussion of aqua regia in 
Lesson XII) but will not reduce sulfuric acid (see the discussion 
of the preparation of hydrogen chloride and hydrochloric acid in 
Lesson IX). Because of the action of hydriodic and hydro- 
bromic acids on sulfuric acid, the latter cannot be used in pre- 
paring them (pages 270 and 274). 

Fluorine. Because of its great activity, all early attempts 
to prepare fluorine failed. It is so active that it decomposes 
water; electrolysis of water solutions of its salts gives oxygen at 
the anode instead of fluorine. In 1886 Moissan prepared it by 
the electrolysis of a solution of potassium hydrogen fluoride 
(KHF 2 ) in liquid hydrogen fluoride. A platinum or copper tube 
must be used, as the element decomposes glass. (See Fig. 104, 
page 267, and page 265.) The properties of fluorine will be under- 
stood from a study of the previous sections of this lesson. It 
unites with most elements, giving light, but the action with gold, 
platinum, or copper is superficial. It does not combine with 
oxygen. 

Hydrofluoric acid. Hydrogen fluoride, or hydrofluoric acid, 
(H 2 F 2 ) is prepared from calcium fluoride and sulfuric acid: 

CaF 2 +H 2 S0 4 ->CaS0 4 +H 2 F 2 

This action is completed because the acid is formed as a gas 
(page 267). Hydrofluoric acid is a weak acid, its water-insoluble 
salts being soluble in the strong acids. It acts vigorously upon 
organic matter and will produce a very painful and slow-healing 
sore when it comes in contact with the flesh. Glass is decomposed 
by it. Considering glass as partly composed of silicon dioxide 
(Si0 2 ) the reaction is represented thus: 

Si0 2 +2H 2 F 2 ->SiF 4 +2H 2 

The silicon tetrafluoride (SiF4) is a gas, which accounts for the 
completeness of the reaction. The acid must be kept in wax 
bottles. It is used in etching designs upon glass and china (page 
268) and in decomposing silicates for analysis. 



92 ELEMENTARY CHEMISTRY 

Chlorine. Chlorine and its compounds have been studied in 
Lesson IX. The first part of that lesson should be reviewed as a 
part of the work at this time. 

Bromine. Bromine is found almost entirely in the form of 
sodium bromide (NaBr) and magnesium bromide (MgBr 2 ), which 
are found in many salt waters. The word "bromine" means a 
stench, referring to the disagreeable odor of the element. It is 
one of the two liquid elements, the other being mercury. It can 
be prepared from its salts by the action of manganese dioxide 
(Mn0 2 ) and sulfuric acid. The complete equation is 

2NaBr+2H 2 S0 4 +Mn02->Na2S0 4 +MnS04+2H 2 0+Br 2 

Commercially, bromine is obtained by the electrolysis of salt 
water. If chlorine is set free in the process, it acts on the bro- 
mide to form bromine as follows: 

2NaBr+Cl 2 ->2NaCl+Br 2 

This is because chlorine is more active than bromine and will 
liberate it from its compounds, just as zinc will liberate hydrogen 
or copper from their compounds. Bromine acts much like chlorine 
but is less active. It is used in preparing bromides and certain 
organic compounds as drugs and dyes. 

Hydrobromic acid. Hydrogen bromide, or hydrobromic acid, 
cannot be prepared pure, as hydrochloric acid can, by the action 
of sulfuric acid on its salts. This is because the sulfuric acid 
oxidizes the hydrobromic acid which is first formed: 

2NaBr+H 2 S0 4 ->Na 2 S0 4 -h2HBr 

Hydrobromic acid is a better reducing agent than hydrochloric 
acid so it reduces the sulfuric acid, forming bromine (page 270): 

H 2 S0 4 +2HBr->2H 2 0+S0 2 +Br 2 

Hydrobromic acid can be made pure by two methods: (1) by 
the action of water, bromine, and red phosphorus as described 
on page 271; (2) by the action of hydrogen sulfide on bromine 
water as described for the preparation of hydriodic acid on page 
274. Both these methods are suitable for preparing both hydro- 
bromic and hydriodic acids. Except that it is less stable and a 
better reducing agent, the properties of hydrobromic acid are 
similar to those of hydrochloric acid. 



ELEMENTARY CHEMISTRY 93 

Iodine. The element iodine is obtained indirectly from sea 
water, either from the ashes of sea weeds or from the residues 
from which Chili saltpeter has been separated. The methods 
given for preparing chlorine and bromine by using manganese 
dioxide and sulfuric acid can be used here (page 269). Chlorine 
will liberate it from sodium iodide: 

Cl 2 +2NaI-*2NaCl+I 2 

Commercially, it is now obtained from sodium iodate (NalCy, 
which occurs in crude Chili saltpeter. It is liberated by the 
action of sulfites of sodium (page 273). 

Iodine forms purplish-black crystals which give a brilliant 
purple vapor. Its odor is irritating but less so than those of 
bromine and chlorine. It can be set free from its salts by either 
bromine or chlorine: 

2NaI+Br 2 ->2NaBr+I 2 

2NaI+Cl 2 ->2NaCl+I 2 

It gives a bright blue color with starch solution, which serves 
as a test either for starch or for iodine as occasion requires. 
Tincture of iodine, a solution of iodine in alcohol, is used as a 
counter irritant in medicine. It is used to prepare iodides, dyes, 
and drugs. Iodoform (CHI 3 ) is used as an antiseptic. 

Hydriodic acid. This acid must be prepared by the methods 
given previously for the preparation of hydrobromic acid (pages 
270 and 274). It is quite unstable and reacts with the oxygen 
of the air to free iodine: 

4HI+0 2 ->2H 2 0+2I 2 

This causes the solution of hydriodic acid to turn brown on 
standing. It is a very strong reducing agent because of the ease 
with which it is decomposed into its elements. Its reactions in 
reducing sulfuric acid are represented as follows: 

H 2 S0 4 ->H 2 0+S0 2 +[0] 
2HI+[0]-^H 2 0+I 2 
S0 2 +4HI->S+2I 2 +2H 2 
S+2HI->H 2 S+I 2 

Sulfur dioxide, sulfur, and hydrogen sulfide are all among the 
products formed in the action. 



94 ELEMENTARY CHEMISTRY 

Salts of the halogens. The salts of the halogen acids are 
generally soluble in water. Exceptions are the chlorides, bro- 
mides and iodides of lead, silver, and mercurous mercury. Silver 
salts are used in photography because of their action toward 
light. Bromides and iodides of potassium are used in medicine. 
Sodium chloride is common table salt. Fluorite, fluor spar, or cal- 
cium fluoride (CaF 2 ), is the most common fluoride. 

Oxygen Compounds. These compounds are not numerous, as 
the halogens have little tendency to combine with oxygen, only 
chlorine and iodine doing so. Chlorine forms three oxides, having 
the formulas, C1 2 0, C1 2 7 , and C10 2 , and four oxygen acids, 
HCIO, HC10 2 , HCIO3, and HC10 4 . In these compounds the 
valence of chlorine varies on the odd numbers from one to seven. 
These compounds are very unstable and are not important in 
themselves. The salts of hypochlorous acid (HCIO) and of chloric 
acid (HCIO3) are of some importance. Hypochlorites are good 
oxidizing agents and are used for bleaching and disinfecting pur- 
poses, as in the purification of water. The chlorates give up 
oxygen easily and are used in the manufacture of explosives and 
fireworks and in preparing oxygen. Hypochlorites and chlorates 
can be prepared from chlorine and the hydroxide of an alkali 
metal and also by electrolytic methods (page 276). 

EXERCISES 

1. Name the halogen elements in the order of increasing atomic 
weights. 

2. What does halogen mean? 

3. Compare the physical properties of the halogen elements. 

4. Write a full discussion of the chemical conduct of the halogens as a 
family. 

5. In what sense are they oxidizing agents? Give the order of their 
oxidizing power. 

6. What kinds of reagents are their hydrogen compounds? Why? 
Name them in order from best to poorest. 

7. How is fluorine made? 

8. Give the properties of fluorine. 

9. Give the method for making hydrofluoric acid. Give the equation. 

10. Give the properties of hydrofluoric acid. 

11. Describe the method for etching glass. 

12. How is bromine made? Give the equations. 

13. State the properties and chemical conduct of bromine. 

14. How is hydrobromic acid made? 



ELEMENTARY CHEMISTRY 95 

15. Write equations for the action of KBr upon H2SO4. 

16. Give the properties, physical and chemical, of hydrobromic acid. 

17. How is iodine made? 

18. State all its properties. 

19. Give the sources of all the halogen elements. 

20. How is hydriodic acid made? 

21. Give the properties of hydriodic acid. 

22. Write equations for the action of H 2 S0 4 and Nal. 

23. Name the water-insoluble halogen salts. 

24. Give the method for preparing hypochlorites and their uses. 

25. Give the methods for preparing chlorates and their uses. 

26. Answer questions 1, 2, 3, 6, 8, 9, 12, and 13, pages 277 and 278. 

27. Solve problems 15, 16, 17, and 18, page 278. 

SEND EXERCISES FOR LESSONS XI=XV TO THE SCHOOL 



LESSON XVI 
MOLECULAR WEIGHTS; ATOMIC WEIGHTS 

Assignment: Chapter XXIII, McPherson and Henderson. 

Chemical laws. Any attempt to find the atomic weights of 
the elements must be based upon the laws of chemical action. 
These have been studied in previous lessons. They are the laws 
of conservation of mass (Lesson I), of definite composition 
(Lesson II), of multiple proportion (Lesson IV), and of combining 
weights (Lesson V). From these laws the theory of the atom was 
deduced. These laws should be reviewed before proceeding 
farther with this lesson. It is recalled that to every element a 
number (determined by experiment) may be assigned which 
represents its combining, or equivalent, weight with every other 
element. The number selected will depend upon the element 
taken as a basis of comparison and the number given to it for a 
standard value. 

Definition of equivalent weight. The natural selection for a 
standard would be 1 gm. of hydrogen, since hydrogen is the 
lightest element. Llydrogen, however, combines with a limited 
number of elements. Oxygen, on the other hand, combines with 
nearly all the elements. For convenience, it is usual to make a 
comparison between the weights of the different elements and a 
weight of oxygen with which they combine. 



96 ELEMENTARY CHEMISTRY 

What number shall we assign to oxygen? If we use 1 gm. 
for oxygen, several elements will have combining, or equivalent, 
weights less than 1, which would be inconvenient. We might 
select 10, 12, 16, 50, or 100 as the weight of oxygen to use as a 
standard. Any of these numbers could be used and by an analy- 
sis of compounds of oxygen numbers could be found showing the 
weights of each which would combine with 10, or 12, or 16 gms. 
of oxygen. 

To have a logical foundation for our standard, it seems best 
to take water, a compound of oxygen and hydrogen, which has 
been analyzed with the highest degree of accuracy. This gives 
a comparison between hydrogen, the natural standard, and 
oxygen, the convenient standard. When the composition of 
water is determined, we find 1 gm. of hydrogen combining with 
7.94 gms. of oxygen or, to use a whole number of oxygen, 1.008 
gms. of hydrogen combines with 8 gms. of oxygen. These 
numbers are directly determined from experiment. It is best, 
then, to use 8 gms. of oxygen as the standard of equivalent, or 
combining, weights rather than any other number. The equiva- 
lent weight of an element is defined as that weight of an element in 
grams which combines with, or is equivalent to, 8 gms. of oxygen. 

Equivalents not atomic weights. The simplest thing would 
be to call the equivalent weights the atomic weights. This is 
apparently impossible, when the facts expressed by the law of 
multiple proportion are considered. Many elements are found to 
combine with several different quantities of oxygen; carbon with 
two, copper with two, iron with three, and nitrogen with five are 
examples. In these cases we are unable to decide which value 
truly represents the weight of the atom relative to the atom of 
oxygen. 

Our problem must be attacked from another angle. The 
determination of the equivalent weight is, however, the first step 
in finding the atomic weight of any element. The equivalent 
weight can be determined with a much greater degree of accu- 
racy than is possible with the other values to be found and serves 
to correct the final result. 

Molecular weights. It is possible to decide which of the 
several equivalent weights of an element, or what multiple of an 



ELEMENTARY CHEMISTRY 97 

equivalent weight, represents the true atomic weight if we can 
learn the molecular weights of the compounds of the element and 
the number of atoms in these compoimds. 

Our knowledge of molecules is largely obtained from the 
study of matter in the gaseous state. It is necessary, then, to 
review the gas laws of Boyle and Gay-Lussac and the kinetic 
theory explaining them (see Lesson III in this text and pages 60 
and 61 in McPherson and Henderson). From the kinetic theory 
the hypothesis of Avogadro was derived. This states that under 
like conditions of temperature and pressure, equal volumes of all 
gases contain the same number of molecules. Accepting this as 
true, and it has had a very thorough testing out in the develop- 
ment of the science, we have a method for finding the molecular 
weights of gaseous compounds, by comparing the weights of 
equal volumes. 

Example. One liter of oxygen weighs 1.43 gms., 1 liter of carbon 
dioxide weighs 1.98 gms.; all the molecules in 1 liter of oxygen weigh 1.43 
gms.; all in 1 liter of carbon dioxide weigh 1.98 gms. The weights of each 
molecule of the two gases are to each other as 1.43 is to 1.98. The molecular 
weight of carbon dioxide equals 1.98 -r 1.43 Xthe molecular weight of 
oxygen. 

Molecular weight of oxygen. The number to be used as the 
molecular weight of oxygen must now be decided. It will depend 
upon the number selected as the atomic weight of oxygen, which 
may or may not be 8 as we shall see, and upon the number of 
atoms in a molecule of oxygen. Again it will be necessary to 
study the composition of water. This time its volume composi- 
tion to form water vapor should be recalled (pages 78-79). 

Qay=Lussac , s law of combining volumes of gases. Other 
gases combining to form a gaseous product show a similar simple 
ratio between the volumes involved. Two volumes of hydrogen 
and 1 volume of oxygen form 2 volumes of water vapor, 3 of 
hydrogen and 1 of nitrogen form 2 of ammonia, and 1 of hydro- 
gen and 1 of chlorine form 2 of hydrogen chloride. The law of 
combining volumes of gases states that when gases combine to 
form gaseous products, there is a small whole-number ratio between 
the volumes of the reacting substances and the volume of the 
product. 



98 ELEMENTARY CHEMISTRY 

Formula of oxygen and water. Applying Avogadro's hypoth- 
esis to the volume composition of water and representing the 
equal volumes by squares, we have 



H 




H 


+ 






water vapor 



If equal volumes contain equal numbers of molecules, we have 
twice as many molecules of water vapor as we have molecules 
of oxygen. Every molecule of water vapor must contain at least 
one atom of oxygen. A complete study of other oxygen com- 
pounds would reveal no compound of oxygen containing less 
volume of oxygen than water. By definition of an atom, this 
quantity of oxygen must represent one atom of oxygen. There- 
fore, having one atom in each molecule, we have twice as many 
atoms of oxygen as we have molecules of oxygen. Oxygen mole- 
cules must, then, contain two atoms each and the formula for 
oxygen is 2 . 

If an examination of other hydrogen compounds should show 
none with less volume of hydrogen than water vapor, then each 
molecule of water vapor would contain one atom of hydrogen, 
and we would have the same number of atoms of hydrogen as 
we have molecules of hydrogen, and each molecule of hydrogen 
would contain one atom. The formula for water vapor would be 
taken as HO, and we could conveniently use 8 as the atomic 
weight of oxygen and 1.008 for that of hydrogen. 

However, we would not examine hydrogen compounds long 
before studying hydrogen chloride. Its volume composition is 
represented as follows: 



H 


+ CI 



hydrogen chloride 



Hydrogen chloride must contain at least one atom of hydrogen 
per molecule. A study of other hydrogen compounds shows none 
containing less hydrogen, therefore, hydrogen chloride does con- 
tain one atom of hydrogen per molecule. But this gives us twice 
as many atoms of hydrogen as we have molecules of hydrogen. 
Therefore, there must be two atoms in each molecule of hydrogen. 
This being the case, we will have to revise our conclusion 
about water vapor, in which we found hydrogen indicated to have 



ELEMENTARY CHEMISTRY 99 

one atom per molecule, since all hydrogen is exactly like all other 
hydrogen, there being no allotropic forms of hydrogen. If hydro- 
gen contains two atoms in each molecule, in the two volumes 
of hydrogen used to form two volumes of water vapor we will 
have twice as many atoms of hydrogen as molecules of hydrogen 
and, therefore, twice as many atoms of hydrogen as molecules 
of water vapor, since the water vapor and the hydrogen molecules 
are equal in number. Each molecule of water vapor must, there- 
fore, contain two atoms of hydrogen. The formula for water 
vapor must be H 2 and not HO. 

Now, if we try to fit the experimentally determined weight 
ratio of hydrogen and oxygen to this formula, 1.008 gms. of 
hydrogen stands for two atoms and 8 gms. of oxygen for one 
atom. A single atom of hydrogen would have an atomic weight 
of little over one-half. This is an inconvenient value to use and, 
to avoid it, chemists have doubled the values in the above ratio, 
taking 16 to represent the weight of one atom of oxygen and 2.016 
for the two atoms of hydrogen. In this way 16 becomes the 
atomic weight of oxygen and 1.008, the atomic weight of hydro- 
gen. With two atoms in a molecule of oxygen, 32 is the molecu- 
lar weight of oxygen. 

Gram=molecular volume. Referring again to the example of 
carbon dioxide and recalling that 1 liter of oxygen weighs 1.43 
gms. and 1 liter of carbon dioxide weighs 1.98 gms., the molecular 
weight of carbon dioxide would be found by the proportion 

1.43 : 1.98 : : 32 : X 
1.98X32 
1.43 
But 

32^ 1.43 = 22.4 

Therefore, in any case the weight of a liter of the gas multiplied 
by 22.4 equals the molecular weight of the gas; for oxygen: 

1.43X22.4 = 32 

In other words, 32 grams of oxygen would occupy 22.4 liters 
under standard conditions. The molecular weight of any other 
gaseous compound in grams will occupy this same volume (22.4 



100 ELEMENTARY CHEMISTRY 

liters) tinder standard conditions. This is called the gram-moleadar 
volume. 

Finding the atomic weight. The second step in finding the 
atomic weight of an element is to multiply the weights of 1 liter 
of all its gaseous compounds by 22.4 to find their molecular 
weights. Third, the per cent of the element whose atomic weight 
is to be found must be determined in all the above compounds. 
Fourth, multiply the molecular weight by the per cent of the 
element in the compound, and we have a value representing the 
weight of all the atoms of that element which the compound con- 
tains. It will be found that these values group themselves around 
certain points, the larger ones being approximate multiples of the 
smallest group. This means that those compounds containing the 
small amounts of the element contain one atom, since we find no 
compound to contain less. This smallest number does not agree 
exactly with other values in its group because of experimental 
errors in finding the weights of one liter of the several gases. 
The larger values show two, three, or four atoms, as the case 
may be, to be contained by the compounds from which they are 
derived. Fifth, the exact atomic weight of the element is found 
by multiplying the equivalent weight, previously determined, by 
some small whole number which will bring it into approximate 
agreement with the above approximate atomic weight. The prod- 
uct thus obtained is the exact atomic weight. See the calculation 
of the atomic weight of nitrogen as worked out on pages 286 
and 287. 

Other methods of determining molecular weights. So far we 
have considered only those compounds which are gases. Other 
methods are useful in studying compounds which are not gases, 
but they are of limited application. From Raoult's laws (page 
285), the molecular weight of substances which dissolve without 
ionic dissociation may be found by finding the lowering of the 
freezing point, or the rise in the boiling point, produced by the 
dissolved substance. 

Molecular formula. In Lesson V a method for finding the 
simplest percentage formula was given. It is now possible to find 
the correct molecular formula for a compound, if we know its 
molecular weight. 



ELEMENTARY CHEMISTRY 101 

Example. The molecular weight of a compound is 78, it contains 7.7 
per cent hydrogen and 92.3 per cent carbon. Find the molecular formula. 

78X0.923 = 71.99, part of molecular weight due to carbon 
78X0.077 = 6.006, part of molecular weight due to hydrogen 

These weights divided by the atomic weights of carbon and of hydrogen 
will give the number of atoms of each element in the compound. 

71.99 -T- 12 = 5.999, or 6 atoms C. 
6.006-^1.008 = 5.958, or 6 atoms H. 

The values do not always equal whole numbers because of errors in the 
experimental work. The molecular formula is C 6 H 6 . 

Molecular equations. Reactions in which elementary gases 

are used should always be written to show the molecular formula 

of the element. If this is done the calculation of quantities can 

be very much simplified by making use of Avogadro's hypothesis. 

Examples. The equation 

2H 2 +0 2 ->2H 2 

shows that two molecules of hydrogen combine with one molecule of oxygen. 
If 10 liters of hydrogen are used, it is clear that 5 liters of oxygen will be 
required to complete the reaction. Gases are usually measured by volume, 
so the volume result is all that is necessary. 

Propane is C 3 H 8 . It burns to water and carbon dioxide, as do all carbon 
and hydrogen compounds: 

C 3 H 8 +50 2 -»3C0 2 +4H 2 

Therefore, 10 liters of propane would require 50 liters of oxygen for its com- 
bustion. 

Ethyl alcohol is C 2 H 5 OH. Its combustion is thus shown: 

C 2 H 5 OH +30 2 ->2C0 2 +3H 2 

Ten liters of alcohol vapor would require 30 liters of oxygen. 

Ethane is C 2 H 6 . In writing the equation for its combustion, we first get 

C 2 H 6 +0 2 ->2C0 2 +3H 2 

We have seven oxygen atoms on the right side. The equation can only be 
balanced by taking the least common multiple of seven and two, the number 
of atoms in a molecule of oxygen. That will multiply 2 by seven and the 
other factors in the equation by two. We now have 

2C 2 H 6 +70 2 ->4C0 2 +6H 2 

Ten liters of ethane would therefore require 10 -f- 2X7, or 35 liters of oxygen. 

EXERCISES* 

1. Define equivalent weight. 

2. State Avogadro's hypothesis. 



*Prepare this set of Exercises and hold it until those for Lessons XVII, XVIII, XIX, 
and XX are also prepared, and then send all five seta to the School. 



102 ELEMENTARY .CHEMISTRY 

3. State Gay-Lussac's law for combining volumes. 

4. Show that oxygen contains two atoms per molecule. 

5. Show that chlorine contains two atoms per molecule. 

6. Show that the formula for water should be taken as H2O. 

7. Define gram-molecular volume. What is its value? 

8. Solve problem 1, page 289. 

9. Solve problems 2 and 3, page 289. 

10. Find the molecular formulas of the compounds from the following 
data: 

(a) mol. wt.=46, % C = 52.17, % H = 13.04, % 0=34.78 

(b) mol. wt.=58, % C = 82.76, % H = 17.24 

(c) mol. wt.=60, % C =40.00, % H = 6.67, % 0=53.33 

11. What volume of the first substance in the equations below will 
combine exactly with 50 liters of the second? Balance the equations: 
2 +CH 3 OH->C0 2 +H 2 0, P 4 +C1 2 ->PC1 6 , HCN+0 2 ->H 2 0+C0 2 +N 2 , 
H 2 S+0 2 ->S0 2 +H 2 0, C 6 H 6 +0 2 -^C0 2 +H 2 0, Hg+0 3 ->HgO 

12. Solve problems 4, 5, 7, 8, 10, and 11, pages 289 and 290. 

13. When ozone acts, it first forms a molecule of oxygen and an atom 
of nascent oxygen. Assuming that 1 liter of ozone oxidizes some substance, 
not acted upon by molecular oxygen, to form a product not a gas, what 
volume of gas is left after the action? 

LESSON XVII 
COMPOUNDS OF CARBON 

Assignment: Chapters XXIV and XXV and Review of Chapter 
X, McPherson and Henderson 

Organic Chemistry. Prior to 1827 chemists believed that 
chemical compounds belonged to two distinct classes. One class 
included all compounds that were formed in living bodies. Matter 
in the bodies of plants and animals is organized into cells and 
tissues, each of which has its function to perform in the life of 
the plant or animal. Because of this organization of matter, the 
compounds found in plants and animals or derived from them 
were called organic compounds. 

The other class included the compounds found in the non- 
living, or mineral, world. The matter in these substances is not 
organized, and they are called inorganic compounds. It was 
believed that some kind of vital force was necessary to make 
the organic compounds, that they could not be made except in 
the living body. This belief is known now to be unfounded. 



ELEMENTARY CHEMISTRY 103 

In 1827 Wohler made urea, a typical organic compound, from 
ammonium cyanate, a compound which can be made from the 
inorganic elements, carbon, hydrogen, oxygen, and nitrogen. 
Since then thousands of organic compounds have been made in 
factories and laboratories. 

It is still customary, for convenience, to retain . the names 
"organic" and "inorganic" as divisions of chemistry, though there 
is no fundamental distinction between them. Organic chemistry 
must, however, be given a different definition. It may be defined 
as the chemistry of the compounds of carbon. Some carbon 
compounds, like the oxides and the carbonates, are more like 
inorganic substances and are studied best in connection with 
other inorganic substances. With this in view, we call organic 
chemistry the study of the compounds of carbon and hydrogen and 
their derivatives. In an elementary course only a few of the 
many organic compounds can be mentioned. 

Carbon monoxide. Carbon and carbon dioxide (C0 2 ) have 
been studied in Lesson VI. Carbon monoxide (CO) is another 
oxide of carbon. It is found in volcanic gases; it may be pre- 
pared by reducing carbon dioxide: 

C0 2 +C->2CO 

This occurs in hard-coal stoves. The blue flame observed at 
times is due to the burning of carbon monoxide above the hot 
coals. The decomposition of formic acid (HCH0 2 ) or oxalic acid 
(H2C2O4) is used for its laboratory preparation (pages 292 and 293). 

It is an odorless gas and an active poison. Deaths frequently 
result from the partial combustion of carbon compounds, as the 
incomplete burning of gasoline in a closed garage. Carbon monox- 
ide is a good reducing agent because of its tendency to combine 
with oxygen. It is the real reducing agent in converting iron 
oxide to iron in the blast furnace. Mixed with hydrogen, it con- 
stitutes water gas, which is used for fuel. 

Carbonates. Carbon dioxide is the acid anhydride of carbonic 
acid (H2CO3). When carbon dioxide is passed into water, the 
solution has the properties of a weak acid. It will neutralize 
bases to form salts, known as carbonates. The acid itself is so 
very unstable that it cannot be isolated. The reaction of water 



104 ELEMENTARY CHEMISTRY 

and carbon dioxide is a reversible one, equilibrium being reached 
when only a small amount of carbonic acid is formed. 

The salts of this acid are quite common. Limestone, shells, 
and marble are calcium carbonate (CaC0 3 ). Washing soda is 
sodium carbonate (Na 2 C0 3 ), and baking soda is sodium acid car- 
bonate (NaHC0 3 ). The carbonates, except those of the alkali 
metals, are insoluble in water. They are acted upon easily by 
acids to liberate carbon dioxide. Calcium carbonate will dissolve 
in a solution of carbonic acid to form calcium acid carbonate 
(Ca(HC0 3 ) 2 ). 

CaC0 3 +H 2 C0 3 ->Ca(HC0 3 ) 2 

This happens in the formation of some kinds of hard water, when 
the water passes over limestone deposits. The acid carbonate can 
be decomposed by heating and the water softened in this way. 

Ca(HC0 3 ) 2 ->CaC0 3 +H 2 0+C0 2 

Cyanogen and hydrocyanic acid. Carbon and nitrogen unite 
at high temperatures to form cyanogen (C 2 N 2 ). This is a very 
poisonous gas. Hydrogen, carbon, and nitrogen form a compound 
called hydrogen cyanide, or hydrocyanic acid (HCN). This is fre- 
quently called prussic acid, because of its relation to prussian blue 
(Fe 4 (Fe(CN) 6 ) 3 ). It is a weak acid. Both the acid and its salts 
are very poisonous. Compounds which can form the cyanide ion 
(CN~) are very poisonous, while those that contain the cyanogen 
radical as a part of some more complex ion are not. The acid and 
its salts are used as bug poisons. Solutions of sodium and potas- 
sium cyanide are used in electroplating and to dissolve gold and 
silver in extracting these metals from their ores. 

Hydrocarbons. Compounds of carbon and hydrogen are 
called hydrocarbons. These compounds are divided into two 
classes called aliphatic and aromatic. The paraffin, or methane, 
series is the principal series of the aliphatic hydrocarbons. The 
term "aliphatic" is taken from the fact that some of the higher 
members of the series are related to the common fats. The 
smaller-molecular-weight members of the methane series are given 
on page 297. Compounds of the same elements differing in com- 
position by a common difference, or common group of atoms, as 
CH 2 in this case, are said to form an homologous series. 



ELEMENTARY CHEMISTRY 105 

Acetylene (C 2 H 2 ) is the first member of another homologous 
series. It has less hydrogen per carbon atom and is said to be 
unsaturated. The aromatic hydrocarbons take this name because 
most of the first known compounds in this class had a pleasant 
odor. The benzene series is the most important of this class. 
Benzene (C 6 H 6 ) and toluene (C 7 H 8 ) are its simplest and most com- 
mon members. They are both obtained from coal tar and coal gas. 

Petroleum. Rock oil, mineral oil, or petroleum, is chiefly 
composed of paraffin hydrocarbons and is the principal source of 
them. The crude oil may be used as a fuel, but it is usually 
refined by distillation. Fractions between different boiling points 
are taken. In the early days it was desirable to get as much 
kerosene (150° to 300°) as possible. Laws had to be passed 
requiring a product with a certain flash point. The temptation 
was to put into the kerosene too much of a lower-boiling fraction. 
If, on heating the oil, its vapor gave a flash of flame with a 
lighted match below a certain temperature, the oil was condemned 
as dangerous to use in lamps. 

The temptation now is to put too much high-boiling fraction 
in the low-boiling fraction. Gasoline (70° to 150°) is now more 
in demand than kerosene. Commercial gasoline has as high a 
boiling point as can be used in the gasoline engine. To obtain 
from petroleum the maximum quantity of hydrocarbons which 
will boil low enough to form gasoline, the so-called cracking of 
oils is. employed; this consists of vaporizing the oils and heating 
the vapor under considerable pressure. One of the methods is 
known as the Burton process. Under proper conditions, benzene 
and toluene are obtained by the process as well as low-boiling 
paraffins (pages 298, 299, and 300). 

Methane (marsh gas). Methane (CH 4 ) is called marsh gas 
because it is formed in marshes by decaying vegetable matter. 
It is also found in mines and therefore is likewise known as fire 
damp. It explodes with oxygen if ignited, the product being 
choke damp (C0 2 ). To protect miners from dangerous explosions 
before mines were electrically lighted, Davy invented the miner's 
safety lamp (page 301). Methane is a constituent of coal gas and 
natural gas. It burns with a pale-blue flame, hence has no 
illuminating power. Methane is made in the laboratory by dis- 



106 ELEMENTARY CHEMISTRY 

tilling a mixture of fused sodium acetate and sodium hydroxide. 
Sodium carbonate and methane are obtained: 

NaC 2 H 3 2 +NaOH->Na 2 C03+CH 4 

Chloroform (CHC1 3 ), iodoform (CHI 3 ), and carbon tetrachloride 
(CCI4) are common halogen derivatives of methane. 

Acetylene. Acetylene (C 2 H 2 ) is made from calcium carbide 
and water: 

CaC 2 +2H 2 0->C 2 H 2 +Ca(OH) 2 

The gas burns with a brilliant white light with plenty of air, 
but otherwise with a very smoky flame. It is used for lighting 
farm houses and was formerly much used in automobile lighting. 
It gives off much heat when burned and forms an intensely hot flame; 
for this reason it is used in the oxyacetylene blow-pipe for cutting 
and welding metal. It is also used to burn carbon out of the 
cylinders of gasoline engines. 

Benzene. Benzene (C 6 H 6 ) and toluene (C 7 H 8 ) are obtained 
from the destructive distillation of coal (page 306) — some in the 
coal tar, more from the coal gas that distills over. Naphthalene 
(CioH 8 ) and anthracene (C14H10) are obtained from the same source. 
Benzene is used as a fuel, as a solvent for fats, and in the manu- 
facture of many useful substances such as dyes, perfumes, drugs, 
flavoring extracts, and explosives. Toluene has the same uses. 
T.N.T., a high explosive, is trinitrotoluene (C7II5N3O6). Naph- 
thalene is used in moth balls and in making indigo. Anthracene 
is used to make alizarin, a red dye. Phenol, or carbolic acid, 
(C 6 H 5 OH) is found in coal tar or can be made from benzene. 
It is a poison and a good disinfectant. Picric acid, an explosive, 
is made from it. Saccharin is prepared from toluene. It is five- 
hundred times as sweet as sugar, but has no food value; it is not 
digested. Other derivatives of benzene are mentioned on page 337. 

Fuels. Wood was doubtless the earliest fuel, and still is an 
important fuel where it is plentiful enough. Coal, in its different 
forms, is the most widely used fuel, but oils and gases are much 
used also. The navy has oil-burning ships, and some railroads in 
Texas and California use oil-burning engines. Gas was first used 
for lighting but now is used largely for fuel as well. There are 
several kinds of fuel gas. Acetylene has already been mentioned. 



ELEMENTARY CHEMISTRY 107 

Others are coal gas, water gas, producer gas, and natural gas. A 
table showing the percentage composition of each is found on page 311. 

Coal gas. The destructive distillation of coal has .been referred 
to when discussing ammonia, coke, and coal tar. Coal was 
first distilled to get coal gas for illumination, the manufacture of 
which is described on pages 306 and 307. The coke made in this 
process is not enough to supply the demand. Much coke is made 
in the bee-hive ovens and in by-product ovens. The former oven 
is very wasteful, as the ammonia, coal tar, and coal gas escape, 
and it is being rapidly replaced by the by-product coke oven. 
These methods are described on page 308. Coal gas burns with a 
luminous flame because of the small amount of acetylene (C 2 H 2 ) 
and ethylene (C 2 H 4 ) which it contains. These gases themselves 
burn with a smoky flame. 

Water gas. Water gas is made by passing superheated steam 
over hot anthracite or coke, carbon monoxide and hydrogen being 
formed (page 308). It is very poisonous, because of the presence 
of carbon monoxide. It has no odor and burns with a non- 
luminous flame; to be used for lighting with an open jet burner, 
it must be enriched with hydrocarbons which will burn with a 
smoky flame. Water gas is enriched with petroleum products by 
passing the gas over hot fire bricks upon which the petroleum is 
sprayed. This cracks the petroleum, forming methane, acetylene, 
and ethylene among other products. Coal gas is sometimes en- 
riched in a similar way. 

Other gases. Producer gas is largely carbon monoxide. It 
can be made from a low grade of coal (page 310). It is used to 
run gas engines and in many metallurgical furnaces. Natural gas 
consists largely of methane; it has moderate luminosity and is 
used as a fuel as well as for lighting (page 310). 

Flames. A flame is the result of the chemical union of two 
gases giving heat enough to produce light. The terms "combust- 
ible" and "supporter of combustion" are relative. We usually say 
that illuminating gas burns in air. Fig. 119, page 311, shows how 
air may be made to burn in an atmosphere of illuminating gas. 
If one of the substances taking part in the action is not a gas, no 
flame results, though light may be given off. Charcoal burns 
without a flame. Liquids and solids, like oil and wood, which 



108 ELEMENTARY CHEMISTRY 

seem to burn with a flame, do so because the heat of the action 
vaporizes them or decomposes them into volatile substances (see 
the illustration with a candle flame, page 314). Note the discus- 
sion of oxidizing and reducing flames, page 318. 

Luminosity of flames. The luminosity of a flame is due to 
the presence of some solid matter which becomes heated to a 
temperature sufficient to give off light. A difference in density 
of the gas in different parts of the flame, caused by decomposi- 
tions, makes the gas visible. In ordinary flames, produced by 
burning carbon compounds, the luminosity is caused by particles 
of unburned carbon heated to high temperature. The carbon is 
formed by the partial combustion of the gas: 

2C 2 H 2 +0 2 ->4C+2H 2 

The gas mantle employs the same principle. The mantle is com- 
posed of thorium and cerium oxides which are stable at high 
temperatures. A non-luminous flame is used to heat the mantle 
to a temperature where it gives off a brilliant white light (page 312). 
Products of combustion of fuels. When fuels are completely 
burned, the products are carbon dioxide and water. These sub- 
stances are not harmful, but burning fuels in rooms without venti- 
lation soon uses up the oxygen necessary for respiration. If 
combustion is incomplete, harmful substances like carbon monoxide 
may form. Complete combustion forms no smoke. Smoke con- 
sists of carbon in a finely divided state, formed from the decompo- 
sition of the gases first formed from the fuel. To prevent smoke 
requires complete combustion of the fuels. Devices for this pur- 
pose provide a sufficient supply of air, thorough mixing of com- 
bustible gases with the air, and a temperature high enough to 
maintain combustion. (See the illustration on page 318.) Various 
fuels differ in their heat-producing, or calorific, value. The values 
for different fuels given on page 320 are obtained by the use of a 
bomb calorimeter (page 321). 

EXERCISES 

1. What is organic chemistry? 

2. Give the methods for preparing carbon monoxide. 

3. State the properties of carbon monoxide. 

4. State the properties of carbonic acid. 



ELEMENTARY CHEMISTRY 109 

5. What is prussic acid? Give its properties. 

6. What are hydrocarbons? Name three series with examples of each. 

7. What is meant by an homologous series? Illustrate. 

8. Discuss the refining of petroleum. 

9. Give the preparation and properties of methane. 

10. Give the preparation and properties of acetylene. 

11. Answer questions 2, 3, 4, 5, 6, 7, 8, 9, and 11, page 304. 

12. Solve problems 12, 13, 14, 15, 16, 17, 18, and 19, pages 304 and 305. 

13. Describe the manufacture of coal gas. 

14. Describe the by-product coke oven. 

15. How is water gas made? 

16. How is producer gas made? 

17. Discuss the composition, properties, and uses of the different fuel 

\ 18. What is a flame? 

19. State the cause of luminosity of a flame. 

20. Explain smoke prevention. 

21. What are meant by oxidizing and reducing flames? 

22. Answer questions 1, 2, 3, 5, 6, 8, and 10, pages 322 and 323. 

23. Solve problems 11 and 12, page 323. 

24. What is meant by the "cracking of oil"? How and why is it car- 
ried out? 

LESSON XVIII 
COMPOUNDS OF CARBON— Continued 

Assignment: Chapters XXVI, XXVII, and XXXIII, McPherson 
and Henderson 

Carbohydrates. We have seen that coal and petroleum are 
important sources of organic compounds. Another large source 
of organic compounds is the class of substances known as carbo- 
hydrates; they include the sugars, starch, cellulose, and such 
substances. The carbohydrates are commonly thought of as com- 
pounds of carbon, hydrogen, and oxygen, with the hydrogen and 
oxygen present in the proper proportions to form water. This 
definition of carbohydrates is not strictly correct {several exceptions 
are known), but it will serve our present purpose. 

Carbohydrates may be divided into classes as: (1) single 
sugars, like dextrose and levulose, both of which have the formula, 
C 6 Hi 2 6 ; (2) double sugars, like sucrose (common sugar), maltose 
(malt sugar), and lactose (milk sugar), all having the same 
molecular formula, C12H22O11; (3) more complex substances, like 



110 ELEMENTARY CHEMISTRY 

starch and cellulose, both of which have a formula of C 6 Hi O 6 
representing their percentage composition, but whose molecular 
weights are unknown. Substances like the sugars, which have 
the same molecular weight but different properties, are called 
isomeric. The difference in properties is due to a different 
arrangement of the atoms in the molecules. 

Sucrose. This sugar is common sugar, familiar in granu- 
lated form. It is obtained from sugar cane and sugar beet 
chiefly. Maple sugar is a variety of sucrose, mostly desired for 
the impurities that flavor it. The sugar from cane and from 
the beet are the same chemical substance (C12H22O11), except for 
small traces of impurities due to their different sources. These 
impurities are too small in amount to affect their properties 
materially, though they may cause slight differences in some 
respects. Popular opinion to the contrary notwithstanding, it 
has been shown that sugar from the beet will make good jelly, 
and pound for pound it is as sweet. The manufacture of sugar 
is discussed on page 325. 

Sucrose does not ferment, but it can be split up into two 
single sugars, dextrose and levulose, by adding a molecule of water: 

C 12H22O11 + H2O— >C eH^Oe + C 6H12O6 

This reaction can be caused by heating the sucrose with a little 
dilute acid, or by the action of certain bacteria. This mixture 
is called invert sugar. Sucrose is decomposed by heat; a sub- 
stance known as caramel is first formed, but the final product is 
carbon. 

Lactose and maltose. Lactose is the sugar found in milk. 
It is added to cow's milk to modify the milk for feeding babies, 
as mother's milk contains a larger percentage of lactose than cow's 
milk. The sugar is added for its food value and not for the 
sweet taste; lactose is not very sweet, and for that reason is 
preferred to sucrose, though both have the same food value. 
When milk sours, lactose undergoes a change caused by bacteria; 
pasteurizing milk kills these bacteria and so keeps the milk sweet. 
When milk sours, lactic acid is formed. 

Maltose is obtained by the hydrolysis of starch. Heating 
starch with a dilute acid, or the action of certain bacteria, first 



ELEMENTARY CHEMISTRY 111 

forms a substance known as dextrin, then maltose. The maltose 
then splits up into two molecules of dextrose. Maltose is formed 
in the process of making fermented liquors. It is also prepared 
in a form mixed with dextrin, which is used instead of lactose to 
modify milk for babies. 

Dextrose. This is found in the juices of fruits and in honey 
together with levulose. It is especially plentiful in grapes and is 
called grape sugar. Commercially it is made from starch by 
heating it with hydrochloric acid. First the starch molecule is 
split up into dextrine, and then this into maltose, which in turn 
splits into dextrose. Most of the commercial product is a mixture 
of dextrose, maltose, and dextrin; it is known as glucose. Dex- 
trose is a good food product but it is not so sweet as sucrose. 

Starch. Starch is found in the seeds and roots of many 
plants; wheat, corn, barley, rice, and potatoes are common 
sources of it. The preparation of starch is described on page 
327. Starch from different sources has a different appearance in 
the shape and size of the grains, so the source can usually be 
detected by the use of the microscope (see the illustrations on 
page 329). 

Cellulose. This substance is the wood fiber of plants; cotton 
and linen are nearly pure cellulose. It is insoluble in most 
reagents but will dissolve in a solution of copper oxide in ammo- 
nium hydroxide. Boiling it with hydrochloric acid forms dextrose. 
Concentrated nitric acid converts it into cellulose nitrates, or 
nitrocellulose, commonly called guncotton, which has many com- 
mercial uses; photographic films, collodion, and celluloid are made 
from guncotton. Mercerized cotton and artificial silk are forms 
of cellulose (page 329). Paper consists mainly of cellulose. 
Cotton rags, linen, and wood pulp are used in making the differ- 
ent grades of paper. The method of manufacture is described on 
page 330. 

Alcohols. Alcohols are a class of organic compounds which 
may be defined as hydroxides of the hydrocarbon radicals. They 
are the simplest oxidation products of the hydrocarbons, though 
they are not made from the hydrocarbons. They are usually 
obtained from natural sources. '* Three important alcohols may be 
considered here; they are methyl (CH 3 OH), ethyl (C 2 H 5 OH), and 



112 ELEMENTARY CHEMISTRY 

glyceryl (C 3 H 5 (OH) 3 ). The latter is the common substance known 
as glycerin. It is related to the hydrocarbon propane, three 
hydrogen atoms being replaced by three hydroxyl groups. 

Methyl alcohol. This alcohol is commonly called wood alco- 
hol. It is one of the products of the destructive distillation of 
wood. It has a lower boiling point than common alcohol and is 
very poisonous. It is used as a solvent and as a fuel. When its 
vapor, mixed with air, is passed over hot copper, oxidation takes 
place with the formation of formaldehyde (CH 2 0). This substance 
is a gas. A solution of it is commonly used as a preservative and 
disinfectant. 

Ethyl alcohol. The common name of ethyl alcohol is grain 
alcohol. It is made by the fermentation of sugars, dextrose 
chiefly, by the action of common yeast (page 333). If starch is 
used in its preparation, it is first converted to dextrose, which 
ferments. Ethyl alcohol is intoxicating and in large amounts is a 
poison; it burns; it is used as a fuel, as a solvent, and in the prep- 
aration of other organic compounds. Ether, (C 2 H 5 ) 2 0, is made 
from it by the action of concentrated sulfuric acid; ether is used 
as an anesthetic in surgical operations. 

For many purposes the alcohol is just as useful if rendered 
unfit for drinking by introducing into it bad smelling and tasting 
or poisonous substances. Methyl alcohol, benzine, and pyridine 
are used for this purpose. This product is called denatured alcohol 
and is free from the heavy government tax. In bread making 
alcohol and carbon dioxide are formed by fermentation caused by 
the yeast used. The carbon dioxide gas expands and escapes 
through the dough, making it light; on baking, the alcohol escapes 
also. 

Glycerin. This alcohol is a by-product of the manufacture 
of soap. It is studied with soap in Chapter XXXIII, page 416. 
Alcohols are not bases but they do react with acids to form water 
and compounds like salts; these compounds are not salts and are 
called esters. The common fats are esters of glycerin and certain 
organic acids of large molecular weight. Both animal fats and 
vegetable oils belong to this class of compounds. When the fats 
are hydrolyzed by boiling with dilute acid or alkali they form 
glycerin and the acid or its alkali salt. 



ELEMENTARY CHEMISTRY 113 

CsH 6 (Ci8HttQ2)8+3NaOH->C8H6(OH)8+3NaCi8H86Q 2 

fat glycerin soap 

The above reaction is typical of what takes place in the formation 
of glycerin and soap. Glycerin is an oily liquid with a sweet taste. 
Its principal use is to prepare glyceryl nitrate, or nitroglycerin. 
It is also used in medicine. 

Explosives. A rapid chemical action forming gases from 
liquids or solids causes an explosion because of the great volume 
change, the greater the volume changed the greater the explosion. 
Many explosives are organic compounds: the nitroglycerin above 
referred to is one; nitrocellulose is another; trinitrotoluene and 
picric acid are others (pages 420 and 421). These substances are 
used in making the smokeless powders. Dynamite is nitroglycer- 
ine absorbed in wood pulp with sodium nitrate added. Trinitro- 
toluene (T.N.T.) is a high explosive much used in the late war. 
Old-fashioned black gunpowder is a mechanical mixture of carbon, 
potassium nitrate, and sulfur, dry, finely ground, and in proper 
proportions to complete the equation 

2KN0 3 +3C+S->K 2 S+3C0 2 +N 2 

Organic acids. The organic acids are oxidation products of 
the alcohols and are related to the hydrocarbons in a similar way. 
The most common series is known as the fatty-acid series because 
several of the higher members are fatty substances. The fatty- 
acid series is related to the paraffin hydrocarbons as follows: 
Methane (CH 4 ), formic acid (HCH0 2 ); ethane (C 2 H 6 ), acetic acid 
(HC 2 H 3 2 ). The common fats are esters of palmitic and stearic 
acids of this series and of oleic acid of another series. The 
glyceryl ester of butyric acid is the constituent of butter that dis- 
tinguishes it from other fats. These acids are monobasic and weak. 

Acetic acid. This acid was one of the earliest acids known. 
It is the sour constituent in vinegar. Commercially, it is obtained 
as a product of the destructive distillation of wood (page 120). 
Acetic acid forms many salts with the metals, as well as esters 
with the alcohols. Lead acetate (Pb(C 2 H 3 2 ) 2 ), known as sugar of 
lead, is used in sugar analysis. Sodium acetate (NaC 2 H 3 2 ) is 
another. Ethyl acetate (C 2 H 6 C 2 H 3 2 ) is an example of an ester 
of the acid. Vinegar is a dilute solution of acetic acid, together 



114 ELEMENTARY CHEMISTRY 

with certain flavoring substances derived from the fruit juices 
from which it is obtained. In making cider vinegar, the sugar of 
the apple juice is first fermented to alcohol by yeast and then 
oxidized to acetic acid by the bacteria found in "mother of vine- 
gar/ ' Distilled vinegar is simply dilute acetic acid made from 
pure dilute alcohol. It leaves no solids on evaporation. Other 
vinegars do, and an analysis of these solids will show the source 
of the vinegar. 

Fats. The common fats and oils are esters of oleic, palmitic, 
and stearic acids with glycerin. These esters are called olein, 
palmitin, and stearin. They have different melting points. The 
various fats are mixtures of these three in different proportions. 
Olive oil and cottonseed oil are largely olein. Lard contains more 
palmitin and stearin. Tallow is largely stearin. Butter is mostly 
composed of these fats, but its characteristic flavor is due to bu- 
tyrin. Butyrin is more readily hydrolyzed than the other fats, 
forming gtycerin and butyric acid. This causes the butter to be- 
come rancid in taste and smell. Oleomargarine has the same 
composition except for the butyrin; it will not become rancid as 
soon as butter. It equals butter in food value but lacks its 
flavor. In butter there are also what are known as vitamines, or 
substances containing "the growth principle. " These substances 
are found in milk, green vegetables, and near the surface of seeds. 
Lard substitutes can be made from oils, as cottonseed oil, by add- 
ing hydrogen. Olein is thus converted into stearin. 

Foods. Food is classified into fats, carbohydrates, proteins, 
mineral matter, and water. Vitamines may be added to this list. 
Fats and carbohydrates have been discussed. They are composed 
of carbon, hydrogen, and oxygen. Proteins are substances con- 
taining nitrogen in addition to the above elements. Albumin, 
casein, and the gluten of wheat are protein-containing substances; 
lean meat and dried beans are also rich in protein. Protein is 
the essential substance in building up the tissues of the body. 
Fats and carbohydrates are oxidized to furnish the heat of the 
body. Observe the composition of various foods as given on 
page 344. 

Soap. Soap is either a sodium or a potassium salt of the 
acids whose esters are found in the common fats; sodium forms a 



ELEMENTARY CHEMISTRY 115 

hard soap and potassium a soft soap. The fats or the vegetable 
oils are treated with sodium or potassium hydroxide. Usually the 
substances are heated together, 'but there is a cold process. The 
reaction is the same as for the preparation of glycerin. The 
commercial process is described on page 417. Soap is soluble in 
water and reacts slightly with it, giving an alkaline reaction. 
Calcium and magnesium compounds form insoluble precipitates 
with soap. Water containing salts of these metals is known as 
hard water, as they do not lather with soap. Soap cleans largely 
because of its ability to form an emulsion (page 387). 

EXERCISES 

1. What are carbohydrates? 

2. Classify carbohydrates. Give examples of each class. 

3. What are isomeric substances? Give an example. 

4. How is sucrose made? What are its sources? 

5. State the chemical behavior of sucrose. 

6. State the source and properties of lactose. 

7. What is the source of maltose? 

8. How is dextrose made? What are its natural sources? 

9. Describe the manufacture of starch. 

10. What is cellulose? What important substances are made from it? 
Give their uses. 

11. How is paper made? 

12. What are alcohols? 

13. Give the name and the formula for three alcohols. 

14. State the preparation and the properties of wood alcohol. 

15. Give the preparation and the properties of grain alcohol. 

16. What is denatured alcohol? 

17. How is glycerin made? Give its uses. 

18. Give the preparation of the common explosives. 

19. What are the fatty acids? Give two examples. 

20. Describe the preparation of acetic acid. What are its properties? 

21. How is vinegar made? How does cider vinegar differ from acetic acid? 

22. What are the fats? Give examples. What is butter? What is 
oleomargarine? 

23. Give the classes of foods and examples of each class. 

24. How is soap made? 

25. Explain the cleansing action of soap. What causes hard water? 

26. Answer questions 5, 7, 8, and 9, page 338. 

27. Answer questions 6, 8, 9, 10, and 11, page 422. 

28. Solve problems 10, 11, 13, 14, and 15, page 338. 



116 ELEMENTARY CHEMISTRY 

LESSON XIX 

THE PHOSPHORUS FAMILY; SILICON AND BORON 
THE PHOSPHORUS FAMILY 
Assignment: Chapter XXVIII, McPherson and Henderson 

The phosphorus family. The phosphorus family includes nitro- 
gen as well as the elements listed at the top of page 346. Their 
properties vary with increasing atomic weight, as in the chlorine 
family. Here, however, chemical activity is not so great and it 
is not to our advantage to compare them in that respect. More 
can be learned about their properties by comparing their tendency 
to increase in metallic properties as atomic weights increase. All 
the elements of this family form two series of compounds. In 
one they have a valence of three, in the other, five. The element 
is more positive, or metallic, in properties when acting in its 
trivalent form. A consideration of each element in these two 
respects will be worth while. 

Nitrogen is decidedly a non-metal. It is sometimes positive 
as in the oxides, but its oxides which are anhydrides form acids. 
The trivalent one, nitrous acid, is rather weak, but still distinctly 
an acid. The acid of pentavalent nitrogen, nitric acid, is a very 
strong acid, showing that the high-valent form is more negative 
and a better acid former. Nitrogen halides are very unstable as 
nitrogen has little tendency to combine positively. 

Phosphorus, with the next higher atomic weight, is a solid, 
has a waxy luster, but is non-metallic in physical properties. In 
chemical action it is sometimes positive, as in combination with 
oxygen and the halogens. These compounds are more stable than 
the corresponding ones for nitrogen. Phosphorus forms no salts 
with the acids. Its oxides are always acid-forming and never 
base-forming anhydrides. Its acids are both weak, but phos- 
phoric acid (H3PO4), derived from the pentoxide (P2O5), is the 
stronger. 

Arsenic shows more positive properties and weaker negative 
properties. It is, however, still strongly negative, that is, non- 
metallic in behavior. It is a black solid, with dull metallic luster. 
Its oxides are anhydrides of weak acids, of which the one con- 



ELEMENTARY CHEMISTRY 117 

taining pentavalent arsenic is the stronger. Arsenic forms a sul- 
fide which can be precipitated as metallic sulfides are precipitated. 

Antimony has a bright metallic luster and appears like a 
metal. Its trioxide (Sb 2 3 ) is equally a basic and an acid anhy- 
dride. The trioxide is insoluble in water, but will dissolve equally 
well in hydrochloric acid, forming a chloride (SbCl 3 ), or in sodium 
hydroxide, forming a sodium salt (Na 3 Sb0 3 ). The hydroxide 
(Sb(OH) 3 ) is equally an acid and a base. Several other metallic 
hydroxides have acid properties as well as basic properties, such 
as those of aluminium, zinc, and lead. Such hydroxides are called 
amphoteric hydroxides (page 361). Antimony exhibits its more 
metallic properties by forming a sulfate. It does not form a 
nitrate. 

Bismuth is decidedly a metal. It not only has metallic luster 
and the other physical properties of a metal, but it acts chemically 
like a metal. All the other elements of the family form com- 
pounds with hydrogen; bismuth does not. Its oxides are but 
slightly acid forming. The trioxide is predominately a basic 
oxide; it reacts with all the common acids to form salts. Bismuth 
nitrate, carbonate, and phosphate are easily formed salts; such 
salts are not known for other members of the family. 

Phosphorus. One of the principal sources of phosphorus is 
the bones of animals, where it occurs as calcium phosphate. 
Calcium phosphate also occurs as a mineral in the form of phos- 
phorite. Phosphorus is a very active element with oxygen and 
for that reason does not occur free in nature. Phosphorus is a 
necessary constituent of plants and must occur in fertile soils in 
soluble form; calcium phosphate is used in the manufacture of 
fertilizers. Phosphorus is made from bone ash or phosphate rock 
by heating with silica (Si0 2 ) and carbon in an electric furnace 
(page 348). 

The element is known in several allotropic forms, of which the 
white or yellow is the common form. The red form is used for 
many purposes. The two forms differ in properties. The first is 
poisonous, easily ignites, and is soluble in carbon disulfide. The 
second is not poisonous, does not ignite so easily, and is not 
soluble in carbon disulfide. Yellow phosphorus combines slowly 
with oxygen at ordinary temperatures giving a pale light in the 



118 ELEMENTARY CHEMISTRY 

dark. It is liable to spontaneous combustion. The change from 
yellow to red phosphorus is a reversible action (page 349). 

Uses of phosphorus. Yellow phosphorus is used as a poison 
for rats. Both forms are used in chemical work. The chief use 
of the element is in the preparation of matches. Matches were 
formerly made from yellow phosphorus, but on account of its 
poisonous effect on the workers, its use is either prohibited or 
taxed out of existence. Instead, a sulfide (P 4 S 3 ) is used. With 
this an ordinary oxidizing agent and a binding material are used. 
Safety matches are red phosphorus and an oxidizing agent on the 
box and antimony sulfide and an oxidizing agent on the match 
tip (page 350). 

Acids of phosphorus. The two oxides of phosphorus are the 
trioxide (P2O3) and the pentoxide (P2O5); both form acids of phos- 
phorus in combination with water: 

P2O3-I-H2O— > (hypophosphorous) 
P 2 03+3H20->2H 3 P03 (phosphorous) 
P2O5+H2O-+2HPO3 (metaphosphoric) 
P 2 5 +2H 2 0->H4P207 (pyrophosphoric) 
P 2 5 +3H 2 0->2H3P04 (orthophosphoric) 

The last is the most important, and its salts are the common 
phosphates. It is usually called phosphoric acid without the pre- 
fix. It is formed when the pentoxide reacts with hot water; with 
cold water, the pentoxide forms metaphosphoric acid. Pyrophos- 
phoric acid is formed from orthophosphoric acid by heating it to 
225°. The salts of the phosphorous acids are not frequently met 
with. 

Arsenic. This element 6ccurs largely in the free state, as 
sulfides, and as the oxide (page 355). It is prepared by the reac- 
tions shown on pages 355 and 356. Its chief use is as an alloy with 
lead, used in making shot. Most of its compounds are poisonous, 
but the element is not. Nitric acid converts arsenic to arsenic 
acid. Hydrochloric acid has no action on it. With nascent 
hydrogen it forms arsine (AsH 3 ) . Marsh's test for arsenic consists 
of the preparation of arsine and a study of the properties by 
which it can be recognized (page 357). Very small amounts of 
arsenic can be detected by this method. Arsenic trioxide (As 2 3 ) 



ELEMENTARY CHEMISTRY 119 

is commonly called white arsenic. It is frequently given as a 
poison and is used in the manufacture of glass, dyes, and insecti- 
cides. The acids of arsenic and their salts are similar in composi- 
tion and properties to those of phosphorus. Paris green is a 
copper salt of complex formula used as an insecticide. Lead 
arsenate is used to spray trees. 

Antimony. The mineral stibnite (Sb 2 S 3 ) is the source of anti- 
mony, which is prepared by heating the sulfide wi£h iron (page 
360). Nitric acid converts antimony to an insoluble oxide; aqua 
regia forms antimony chloride; hydrochloric acid does not act on 
it. It forms oxides and acids similar to those of phosphorus, but 
much weaker. Stibine (SbHs) is formed in Marsh's test, if anti- 
mony salts are present. The two elements give different reactions 
and can easily be distinguished (page 360). Antimony is used in 
alloys. Antimony salts react with water; but the hydrolysis is 
only a partial one. In antimony chloride only two of the chlorine 
ions are replaced by hydroxyl ions (page 362). 

SbCl 3 +2HOH-^Sb(OH) 2 Cl+2HCl 

This forms a basic salt, which, however, loses a molecule of water 
and becomes SbOCl; this is called antimony oxy chloride. 

Bismuth. Bismuth is usually found in the free state and is 
prepared by melting the ore and allowing the liquid to run into 
vessels (page 363). Bismuth forms the nitrate with nitric acid 
and the sulfate with hot concentrated sulfuric acid, but hydro- 
chloric acid does not act with it. Its principal oxide is the 
trioxide (Bi 2 3 ), which is mostly basic in its properties. Bismuth 
salts hydrolyze partially like those of antimony. Bismuth oxy- 
nitrate (BiON0 3 ), called bismuth subnitrate, is used as a 
medicine in certain kinds of stomach trouble. Bismuth is used 
in alloys. 

Alloys. Alloys are substances formed by melting metals 
together. While liquid, a solution is formed. On cooling we 
have what is called a solid solution. Most alloys are of this kind. 
Some are known to be definite compounds of the metals with 
each other. Lead and copper are metals much used in making a 
great variety of alloys. The many varieties of steel are formed 
by making iron alloys with different metals and with carbon. 



120 ELEMENTARY CHEMISTRY 

The melting point of most alloys is below the average melting 
point of the constituents and in many cases below that of any 
of them. 

Antimony alloys expand on solidifying and for that reason 
are suitable for making type. Type metal consists of antimony, 
tin, and lead. Babbitt metal contains the same elements with a 
little copper added. Bismuth alloys have a very low melting 
point; some of them melting below the boiling point of water. 
Wood's metal, which consists of bismuth, lead, tin, and cadmium, 
melts at about 60°. These alloys are used in automatic fire- 
extinguishing apparatus (page 365). Brass is an alloy of copper 
and zinc. Other copper alloys are listed on page 497. Alloy 
steels and their uses are given on page 481. Nickel, tin, zinc, 
and aluminium are also used in different alloys. Solder is an 
alloy of lead and tin. Coin metals are alloys of copper, nickel, 
silver, and gold. 

SILICON AND BORON 
Assignment: Chapter XXIX, McPherson and Henderson 

Silicon. This element belongs in Periodic Group IV with 
carbon and resembles carbon in some respects. The valence of 
silicon is four, and in the silicates we find a large number of 
complex compounds. As silicon dioxide (Si0 2 ) or silicates of the 
metals it is a constituent of a large part of the substances which 
make up the crust of the earth, such as granite, sandstone, clay, 
and shale. Silicon is made by reducing silicon dioxide with 
aluminium or carbon. It is used in the metallurgy of iron as 
ferrosilicon. Silicon forms silicides with several elements. Carbo- 
rundum, or carbon silicide (CSi), is a very hard substance used 
as an abrasive in place of emery. For other properties of silicon 
see page 368. Carborundum is made in an electric furnace 
(page 370). Silicon dioxide occurs in nature in a great variety 
of forms, as quartz, sand, onyx, opal, flint, and agate. Infusorial 
earth is a form of silicon dioxide derived from the skeletons of 
certain microorganisms; it is used in making scouring soaps. 
Quartz is used to make dishes for laboratory use. Silicon dioxide 
is the anhydride of two acids, orthosilicic acid (H^iOJ and 
metasilicic acid (H 2 Si0 3 ); many of the simple silicates are salts 



ELEMENTARY CHEMISTRY 121 

of the latter, as calcium metasilicate (CaSi0 3 ) ; mica (KAlSi0 4 ) is 
a salt of the former. 

A large number of silicates are complex compounds containing 
the acid radical of what are called condensed silicic acids; that is, 
two or more molecules of silicic acid (H 4 Si0 4 ) lose one or more 
molecules of water to form the condensed acid. Thus: 

3H 4 Si0 4 ->H 4 Si 3 8 +4H 2 

Feldspar (KAlSi 3 8 ) is a salt of this acid. Pure clay, or kaolin, is 
Al 2 Si207.2H 2 0. Water glass is sodium metasilicate (Na 2 Si0 3 ). It 
is used to preserve eggs, in waterproofing porous materials, as a 
glue for glass and pottery, to make curtains non-inflammable, and 
in soaps (page 374). 

Glass. Glass is a solid solution of different silicates and 
silicon dioxide. Common glass contains the silicates of sodium 
and calcium. In other glasses, potassium is substituted for sodium 
to give hardness and a higher melting point. Lead is used in 
place of calcium to give a low-melting glass. Lead and barium 
both give a high refractive index. Lead glass is used in making 
optical instruments. Glass is colored with various metallic oxides 
which fuse along with the glass; milky glass has tin oxide or some 
white infusible substance present. Enamels for surfacing metal 
vessels are a kind of opaque glass which contain the oxides of 
boron in place of some of the silicon dioxide, and oxides of zinc, 
lead, or barium in place of some of the calcium. A full descrip- 
tion of the manufacture of glass is given on pages 374-377. 

Boron. Boron occurs in the third periodic group. Other 
members of this group have metallic properties, but boron is a 
non-metallic element. Its oxide (B 2 3 ) is an acid anhydride, 
forming boric acid (H 3 B0 3 ). This acid occurs in nature. Another 
source of boron is borax (Na 2 B 4 7 ). This compound is a salt of 
tetraboric acid, a condensed acid of boric acid. Boric acid can be 
made from borax by treating it with sulfuric acid: 

Na 2 B 4 7 +5H 2 0+H 2 S0 4 ->Na 2 S0 4 +4H 3 B0 3 
Borax is a hydrate with either five or ten molecules of water. It 
is used in making enamels and glazes. Borax fuses, loses its 
water, and forms a glass which will dissolve metallic oxides with 
characteristic colors in many cases. It is used in analysis to 



122 ELEMENTARY CHEMISTRY 

detect these metals. Cobalt gives a deep blue, chromium a 
green, and iron a yellow color. Borax, being a salt of a weak 
acid and a strong base, has an alkaline action toward litmus. It 
is used to soften water, as an antiseptic, and in brazing. 

EXERCISES 

1. Name the elements of the phosphorus family in order of increasing 
atomic weights. 

2. Discuss the variation of their positive and negative properties as 
atomic weight increases. 

3. Compare positive and negative properties of the two different 
valence forms. 

4. Compare the properties of the yellow and the red forms of phos- 
phorus. 

5. Give the composition of the different kinds of matches. 

6. Write equations showing the relation of the oxides of phosphorus to 
the acids. 

7. Name the sources of calcium phosphate. What is its use? 

8. State the action of arsenic, antimony, and bismuth with hydro- 
chloric and nitric acids. 

9. Describe Marsh's test for arsenic. 

10. Give the uses of arsenic, antimony, and bismuth as elements. 

11. Explain the hydrolysis of bismuth nitrate. 

12. What are alloys? Name five alloys, their composition, and uses. 

13. What is silicon used for? Give its sources. 

14. What is carborundum? How is it made? What is its use? 

15. Name a number of sources of silicon dioxide. What are its uses? 

16. Give names and formulas for the acids of silicon. Name a salt of 
each acid. 

17. What is water glass? How is it made? What are its uses? 

18. Describe the manufacture and the blowing of glass. 

19. What is glass? What are the varieties of glass? How is it colored? 

20. How is boric acid made? What is borax? Give its uses. 

21. Answer questions 2, 5, 6, 7, 9, 11, and 12, pages 365 and 366. 

22. Solve problems 8, 15, 16, and 17, page 366. 

23. Answer questions 2, 4, 5, 6, and 7, page 381. 

LESSON XX 

METALS; ALKALI FAMILY 

METALS 

Assignment: Chapter XXXI, McPherson and Henderson 

The metals. The elements have already been referred to as 
acid forming and base forming. From the point of vieiv of their 
chemical reactions, the metals may be defined as those elements which 



ELEMENTARY CHEMISTRY 123 

can become cations of bases. As we saw in the study of antimony, 
the distinction between a metal and a non-metal is not a sharp 
one. This element and several others have both metallic and non- 
metallic properties. In most cases one or the other kind of prop- 
erty predominates. The metals are solids, except mercury, which 
is a liquid. Most of them have a high density, though the alkali 
metals are lighter than water. They are conductors of heat and 
electricity and have a high luster. Most of them combine readily 
with oxygen and other negative elements, their surfaces tarnishing 
if the action is not complete. The chemical properties of the 
metals are in accord with their position in the electrochemical 
series (page 191), a review of which at this time would be profitable. 

Extraction of metals. Most of the metals are found in nature 
in combination, but gold and frequently silver and copper are 
found free. Minerals are inorganic substances found in nature. 
An ore is a mineral from which a useful substance can be extracted. 
The extraction of metals from their ores is called metallurgy. 

Two methods of metallurgy are in use, reduction and elec- 
trolysis. Reduction may be accomplished with carbon, as in the 
case of iron, zinc, and tin, if the ore is an oxide. Many ores 
which are not oxides can be converted into oxides by roasting the 
ore; carbonates form oxides on heating; sulfides must be heated in 
a current of air. More refractory oxides are not reduced by car- 
bon, and in such cases aluminium can be used as the reducing 
agent. Chromium and manganese are made this way. 

If the ore is soluble or can be easily converted into a soluble 
compound or if it fuses readily, it can be decomposed by the 
electric current. The metal deposits at the cathode if it does not 
combine with the water present. Sodium, which would combine 
with water, is made by electrolysis of a fused compound. Elec- 
trolysis is used to refine some metals which are obtained in impure 
form by other methods. 

Not all the electrochemical industries are electrolytic. Many 
of them use electric energy simply as a method of generating heat. 
Such are the preparation of carborundum, graphite, and carbon 
disulfide. Water power is essential for the generation of cheap 
electric energy. For this reason many of our electrochemical 
industries are located at Niagara Falls (illustration page 392). 



124 ELEMENTARY CHEMISTRY 

Preparation of compounds of metals. Compounds of the 
metals are prepared by employing the same types of reaction 
which have already been studied. These are addition, or com- 
bination, of elements, decomposition of a compound, substitution 
of one element for another, and double decomposition. Double 
decomposition is most frequently carried out in solution. 

It is necessary that the conditions be such as to permit the 
reactions to go to completion (page 224). This may happen in 
one of three ways: (1) an insoluble gas may be formed; (2) an 
insoluble solid may form; (3) an undissociated molecule may form. 
Reactions of these types can be employed in preparing metallic 
compounds if the proper materials are selected. Some substances 
are not soluble in water or acids and so cannot be brought into 
double decomposition in the usual way. Many such insoluble 
substances are soluble in melted sodium carbonate and can be 
brought into reaction by fusing with it. Since the carbonate 
formed by double decomposition is insoluble in the sodium car- 
bonate, completion of the reaction is made possible. 

We have, then, altogether, seven ways to make compounds of 
metals: 

1. Combination, Zn+O— »ZnO 

2. Decomposition, CaC0 3 — >CaO+C0 2 

3. Substitution, CuS0 4 -f Zn— »ZnS0 4 +Cu 

4. An insoluble gas, CaC0 3 +2HCl->CaCl 2 +C0 2 +H 2 

5. An insoluble solid, Pb(N0 3 ) 2 +Na 2 S0 4 -*Pb§04+2NaN0 3 

6. An undissociated molecule, Ca 3 (P0 4 ) 2 +6HCl-*3CaCl 2 +2H 3 P0 4 

7. Fusion, PbS0 4 +Na 2 C0 3 ->PbC0 3 +Na 2 S0 4 

The preparation of zinc sulfate by using zinc and sulfuric acid, 
hydrogen gas being set free, is an example of method No. 3 and 
not of No. 4, which is a type of double decomposition first of all, 
while in preparing zinc sulfate, the element hydrogen is replaced 
by zinc (substitution). The example given under method No. 6 is 
that of an undissociated molecule because phosphoric acid is a 
very weak acid and is practically undissociated in the presence of 
the stronger hydrochloric acid. 

Insoluble compounds. In order to apply methods 4 and 5 it 
is necessary to know what substances are insoluble. The solubil- 
ity of gases varies somewhat with conditions, but we have found 
it possible to establish conditions whereby hydrochloric acid, 



ELEMENTARY CHEMISTRY 125 

hydrogen sulfide, ammonia, carbon dioxide, and nitric acid are 
insoluble. They are all prepared according to method 4. 

Insoluble solid compounds of the metals can be made by pre- 
cipitation according to method 5. To carry out such a reaction, 
it is necessary to start with two water-soluble substances. One 
must contain the metal of the compound it is desired to prepare; 
the other, the negative radical. Then, when these solutions are 
brought together, the metal and negative radical will combine and 
precipitate out, if they form an insoluble compound. A table of 
the insoluble salts of the several acid radicals is given on pages 
394 and 395. It should be studied carefully and frequently 
referred to. 

ALKALI METALS 

Assignment: Chapter XXXII and Review of Chapter XV, McPherson 
and Henderson 

Characteristics of family. This family occurs in Group I of 
the periodic table. It consists of lithium, sodium, potassium, 
rubidium and caesium. Note the relation between their atomic 
weights and physical properties (page 396). As in the chlorine 
and phosphorus families, these metals show a marked gradation in 
properties as their atomic weight increases. They are the most 
positive family of elements, that is, the strongest metals. Their 
hydroxides are strong bases. As the atomic weight increases, 
their metallic properties become stronger. Potassium hydroxide is 
a stronger base than sodium hydroxide, and caesium hydroxide is 
the strongest of all bases. They occur at the top of the electro- 
chemical series of the metals, showing their very positive properties. 

Except as they vary in strength of chemical activity, they are 
very similar in properties. They always combine with a valence 
of one. All their important salts are soluble in water, and for 
this reason their compounds are frequently found in sea water or 
in salt beds formed by evaporation. Their activity with oxygen 
and water is so great that they must be kept under an oil, like 
kerosene. Of these elements, only sodium and potassium are 
common enough to demand special study. Sodium was studied in 
Lesson IX, which should be reviewed at this point. 



126 ELEMENTARY CHEMISTRY 

Compounds of sodium. Sodium chloride (NaCl), common 
salt, is the chief source of sodium and sodium compounds. It is 
found in salt water and in salt beds in many localities. It is used 
as table salt and in the preparation of soap, glass, soda, bleaching 
powder, and hydrochloric acid. Most salt contains small amounts 
of calcium and magnesium chlorides and for that reason absorbs 
moisture. Shaker salt is kept dry by mixing with it a little finely 
powdered calcium phosphate. 

Sodium sulfate (Na 2 S04) is formed in making hydrochloric 
acid and is used in making sodium carbonate. Its hydrate 
(Na 2 SO 4 .10H 2 O) is known as Glauber's salt. Sodium sulfate is 
used as a laxative in cattle powders. It is also used in making 
glass. Hypo is used by photographers to dissolve unchanged 
silver salts and is employed in bleaching to absorb excess chlorine. 
It has the formula Na 2 S 2 3 , which is sodium thiosulfate. 

Sodium carbonate (Na 2 C0 3 ) is made in two ways. The older 
method, known as the Leblanc process, was devised during the 
French Revolution (page 403). It is now largely replaced by the 
Solvay process, devised by a Belgian chemist. The reactions 
involved in these two processes are shown on page 402. It will be 
noticed that hydrochloric acid is a by-product of the Leblanc 
process. The sale of this enables the few Leblanc plants to con- 
tinue to operate on a paying basis. Sodium carbonate forms a 
hydrate (Na 2 CO 3 .10H 2 O), which is known as washing soda or sal 
soda. It is used to soften water and in the manufacture of glass 
and soap. It hydrolyzes with an alkaline action. It is used in 
laundry work because the alkaline action cuts grease. 

Sodium hydrogen carbonate (NaHC0 3 ), commonly called 
bicarbonate of soda, is baking soda. It is made in the Solvay 
process or by passing carbon dioxide into a solution of sodium 
carbonate. It is used in baking and in making baking powder. 

Sodium nitrate (NaN0 3 ) is Chili saltpeter. It is used in 
making nitric and sulfuric acids and in the manufacture of fer- 
tilizers. This is the only nitrate found extensively in nature. 
The largest deposits are in Chili. Sodium cyanide (NaCN) is 
another important salt of sodium. It is used in gold mining 
because its solutions dissolve elementary gold. The compound is 
extremely poisonous. Sodium hypochlorite (NaCIO) is an unstable 



ELEMENTARY CHEMISTRY 127 

compound, obtained only in dilute solution. It is prepared by the 
reaction shown on page 406. It is used as a bleaching and antiseptic 
agent. Sodium compounds impart a yellow color to a colorless flame. 

Potassium. Potassium compounds are abundant, but many 
of them are silicates and hence insoluble and very hard to decom- 
pose. For this reason it is so difficult to obtain the potassium in 
usable form that these compounds are not an important source of 
potassium. The common source of potassium salts is the large 
deposits of chloride and sulfate at Stassfurt, Germany (page 407). 
Potassium is a necessary element in plant foods, and it must occur 
in water-soluble compounds which can be absorbed by the roots of 
the plants. Plant ashes are a source of potassium carbonate. 

Potassium salts are used in fertilizers, and during the war 
fertilizer manufacture was seriously handicapped by the lack of 
such salts, formerly imported from Germany. Other sources (page 
408) of potassium salts were developed as much as possible but 
the supply could not nearly meet the demand. Potassium and 
potassium hydroxide are made exactly like sodium and sodium 
hydroxide (pages 173 and 176). Saltpeter, potassium nitrate 
(KN0 3 ), is used in making gunpowder, as a preservative, and in 
medicine. It is a good oxidizing agent. It is now made from 
sodium nitrate, but formerly it was obtained from manure by the 
oxidizing action of bacteria on the organic compounds containing 
potassium and nitrogen. Potassium carbonate (K 2 C0 3 ) can be 
made by the Leblanc process, but not by the Solvay process 
because potassium hydrogen carbonate is too soluble to be pre- 
cipitated. Potassium compounds give a purple color to a colorless 
flame (page 414). 

Ammonium compounds. The ammonium radical (NH 4 ) has 
positive properties, forming a cation. The base (NH 4 OH) was 
studied in Lesson XII. Its salts are usually classified with the 
salts of the alkali metals. The radical is univalent, and the salts 
are generally soluble in water. Ammonium salts are found in the 
soil in small quantities, being formed by the decay of nitrogen 
containing organic matter. Commercially, they are obtained from 
the ammoniacal liquors produced in making coke (page 306). 
Ammonium chloride (NH 4 C1), called sal ammoniac, is used in dry 
cells, in soldering, in medicine, and as a reagent. 



128 ELEMENTARY CHEMISTRY 

EXERCISES 

1. What is a metal, a mineral, an ore? Define metallurgy. 

2. Name the methods for obtaining metals from ores. Illustrate. 

3. What is roasting? How is it done and why? 

4. Give the seven methods for making compounds of the metals. 
Write two equations to illustrate each. 

5. State the solubility of chlorides, sulfates, carbonates, nitrates, and 
phosphates. 

6. Answer questions 1, 2, 3, and 4, page 395. Assign each equation 
written to one of the seven methods for making compounds. 

7. Discuss the characteristics of the alkali metals. 

8. Give chemical name, formula, and use for the following: caustic 
soda, washing soda, sal ammoniac, caustic potash, saltpeter, baking soda, 
borax, hypo, Glauber's salt, Chili saltpeter. 

9. Describe, giving equations, the Leblanc and the Solvay processes 
for making sodium carbonate. 

10. How are potassium hydroxide and potassium carbonate made? 

11. Give the flame tests for sodium and potassium. 

12. Answer questions 2, 4, 5, 9, 11, and 18, pages 414 and 415. 

13. Solve problems 7, 8, and 13, page 415. 

SEND EXERCISES FOR LESSONS XVI-XX TO THE SCHOOL 



LESSON XXI 

THE CALCIUM AND MAGNESIUM FAMILIES 
CALCIUM FAMILY 
Assignment: Chapter XXXIV, McPherson and Henderson 

Characteristics of family. The metals of this family are cal- 
cium, strontium, and barium — sometimes called the alkaline-earth 
metals. They form weaker bases than the metals of the alkali 
family. Many of their compounds are found in the earth's crust 
as carbonates and silicates. Calcium is by far the most abundant. 
None of them occurs free in nature as they are too active, though 
not so active as the alkali metals. They follow the alkali metals 
in the electrochemical series. They act slowly with water at 
ordinary temperatures, forming hydrogen. They burn brilliantly 
in oxygen or the air. With a colorless flame, calcium compounds 
give a yellowish-red, strontium a crimson, and barium a green 
color. Strontium and barium compounds are used in making 
colored fires. 



ELEMENTARY CHEMISTRY 129 

Compounds of calcium. Calcium oxide (CaO), commonly 
called lime, or quicklime, is obtained by heating calcium carbon- 
ate in the form of limestone or marble. It is used to make 
slaked lime by adding water. This forms calcium hydroxide 
(Ca(OH) 2 ). Slaked lime is used extensively in building as a con- 
stituent of mortar and plaster. Quicklime absorbs carbon dioxide 
from the air as well as moisture, so air-slaked lime contains calcium 
carbonate. Lime is used as a soil dressing because it neutralizes 
the organic acids in sour soil, thus producing more favorable 
conditions for bacterial action to convert organic nitrogen into 
forms available for plant food. 

Limewater is a solution of calcium hydroxide. It is used 
some in medicine. Calcium hydroxide is used in making ammonia 
and bleaching powder and to remove hair from hides. Mortar is a 
mixture of calcium hydroxide and sand; plaster contains hair in 
addition. Mortar hardens and sets because it absorbs carbon 
dioxide from the air to form calcium carbonate (page 427). 

Bleaching powder is made by passing chlorine over lime. 
The compound formed is a double salt containing two acid radi- 
cals (CaClOCl), being a salt of hydrochloric (HC1) and hypo- 
chlorous (HOC1) acids. It is sometimes called chloride of lime. 
It is used for bleaching, and as a disinfectant because of the 
chlorine liberated. It is used to purify city water. 

Calcium carbonate occurs in nature in a variety of forms, as 
limestone, marble, pearls, coral, shells, and chalk (page 428). It 
is largely used as a building stone and in the manufacture of glass, 
cement, soda, lime, and carbon dioxide. Pure calcium carbonate 
is precipitated from solutions of sodium carbonate and calcium 
chloride. This is called precipitated chalk and is used in making 
tooth powder. 

Calcium sulfate occurs in nature in the form of gypsum, 
which has the formula CaS0 4 .2H 2 0. Gypsum is used in making 
paper and plaster of Paris. When gypsum is heated to about 
125° C, it loses three-fourths of its water, forming plaster of 
Paris (CaS0 4 )2-H 2 0. This compound absorbs water and hardens 
by again forming gypsum. It is used to make plaster casts for 
broken bones, in stucco work, and as a finishing coat in plastering 
walls. 



130 ELEMENTARY CHEMISTRY 

Calcium carbide is made by heating lime and carbon in an 
electric furnace (page 432). It is used in making acetylene, and 
nitrogen passed over hot calcium carbide forms calcium cyanamide 
(page 433). This latter compound (CaCN 2 ) is a useful fertilizer, 
its value depending upon its reaction with water to form ammonia 
This is one of the methods of converting atmospheric nitroger 
into compounds of nitrogen. 

Hard waters and boiler scale. Hard water is caused bj 
compounds of calcium and magnesium in solution. Temporary 
hardness can be removed by boiling, permanent hardness cannot 
Temporary hardness is caused by the presence of acid carbonates 
of calcium and magnesium. When water, containing carbon diox- 
ide, passes over limestone, the acid dissolves the carbonate, form- 
ing calcium acid carbonate (Ca(HC0 3 ) 2 ). This is soluble, but is 
unstable when heated, and boiling the water decomposes it, form- 
ing calcium carbonate. Calcium hydroxide is used sometimes tc 
neutralize the acid carbonate. (See the equations on page 431.] 
Magnesium compounds behave in a similar way. 

Sulfates and chlorides of these metals are dissolved by water 
As these compounds are stable and are not removed by boiling 
such hardness is called permanent. The water is softened bj 
precipitating the carbonates with sodium carbonate (equation 
page 431). On a commercial scale, water is softened by firsl 
adding calcium hydroxide to remove the temporary hardness anc 
then sufficient sodium carbonate to correct the permanent hard- 
ness and react with the excess of calcium hydroxide that was used 
The amounts of the reagents to be used must be determined b} 
an analysis of the water. 

If hard water is used in boilers without softening, on th< 
evaporation of the water, a hard scale, known as boiler scale 
deposits (page 443). Three different substances are found ir 
boiler scale: (1) calcium sulfate; (2) calcium carbonate; (3) mag- 
nesium salts. Calcium sulfate, sparingly soluble in cold water, is 
less soluble in boiling water; it is, therefore, precipitated. Cal- 
cium carbonate is precipitated by boiling, if the water contains 
calcium acid carbonate. Magnesium salts are mostly present as 
the chloride. They are hydrolyzed by boiling water, forming 
magnesium hydroxide and hydrochloric acid (equation, page 444). 



ELEMENTARY CHEMISTRY 131 

The conditions referred to in the previous paragraph that sul- 
fates and chlorides of these metals are not removed by boiling are 
for dilute solutions, but here the conditions are those of a concen- 
trated solution, which explains the seeming contradiction. 

Magnesium hydroxide, which forms a part of the scale, has a 
sort of cementing effect, causing the particles of the scale to form a 
hard compact mass which adheres to the boiler. This sometimes 
cracks and the acid acts with the iron, weakening it. Explosions 
may be caused by the water coming in contact with the super- 
heated iron. 

"Boiler compounds," used to prevent the formation of boiler 
scale, do not prevent the precipitation of these substances, but do 
have a tendency to hold them in suspension or keep them in a 
loose state, so that they can be flushed out with water or blown 
out with steam. Tannic acid is the principal ingredient of 
"boiler compounds." 

Fertilizers. Many of the elements are essential to plant 
growth. Six are very important: carbon, hydrogen, oxygen, 
nitrogen, phosphorus, and potassium. Plants obtain carbon from 
the carbon dioxide of the air and hydrogen and oxygen from the 
water in the soil. The other elements must be present in the soil 
in water-soluble form so that they can be absorbed with the water 
by the roots of the plants. As nitrogen, phosphorus, and potas- 
sium compounds are withdrawn from the soil by growing crops 
they must be replaced. This is done by adding fertilizers. Fer- 
tilizers are mixtures of compounds of these three elements. 

1. Sources of nitrogen. Many of the sources of nitrogen 
compounds for this purpose have already been referred to, such 
as, Chili saltpeter, ammonium sulfate, calcium nitrate (air salt- 
peter), and calcium cyanamide. Organic nitrogenous matter, as 
dried blood, slaughter-house waste, and manure, is an important 
source. The organic nitrogen is converted into nitrates by oxidiz- 
ing bacteria. 

2, Sources of phosphorus. The source of the phosphorus in 
fertilizers is usually calcium phosphate. Ground bones are 
especially valuable because they also contain nitrogen. Likewise, 
phosphate rock is mined for this purpose in large quantities. 
Calcium phosphate is nearly insoluble in water, so it would not be 



132 ELEMENTARY CHEMISTRY 

available as plant food until acids in the soil had dissolved it. 
This would take too long; therefore calcium phosphate is treated 
with sulfuric acid, forming soluble calcium acid phosphate (equa- 
tion, page 437). The mixture of sulfate and acid phosphate 
resulting is known as superphosphate of lime. Slag from the steel 
works sometimes contains enough phosphate to be used to make 
fertilizer. 

3. Sources of potassium. The chief source of potassium com- 
pounds has been the Stassfurt mines in Germany (page 407). 
These salts are used directly in the fertilizer. During the war 
this source of potassium compounds was shut off, and fertilizers 
were low in percentage of potash. Wood ashes was used as far as 
the limited supply would go. The composition of a fertilizer is 
varied according to the crop to be grown, but always contains 
some of these three kinds of material, nitrogen, phosphorus, and 
potassium. Lime is sometimes added to a soil to neutralize its 
sour condition (page 438). 

MAGNESIUM FAMILY 

Assignment: Chapter XXXV, McPherson and Henderson 

Characteristics of family. The metals of this family are 
members of the same group as the alkaline-earth metals. Mag- 
nesium, zinc, cadmium, and mercury constitute the family. 
Mercury is below hydrogen in the electrochemical series and for 
that reason has many properties similar to those of copper and 
silver. It will be studied with them. Magnesium and zinc 
occupy places well above hydrogen and, therefore, are active 
metals. They form weaker bases than the alkaline-earth metals 
and are less active with oxygen and water. Magnesium will lib- 
erate hydrogen from boiling water, and red-hot zinc will do the 
same from steam. Both liberate hydrogen from the acids 
freely. They are tarnished in the air. Zinc acts upon strong 
bases, liberating hydrogen. 

Magnesium. Magnesium is prepared by the electrolysis of 
magnesium chloride or carnallite (page 408). The element is a 
light metal. It burns with a brilliant white light which is rich in 
rays that affect photographic plates, and in the form of a fine 



ELEMENTARY CHEMISTRY 133 

powder it is used in making flash lights for taking photographs. 
Mixed with an oxidizing agent like potassium chlorate, it is used 
in making rockets to light battlefields at night. Some alloys 
contain magnesium as magnalium; they are used where lightness 
is required. 

Compounds of magnesium. The compounds of magnesium 
resemble those of calcium in many respects. Magnesium oxide 
(MgO), a soft bulky powder, is called magnesia. It is unchanged 
at high temperatures and is used in making crucibles and lining 
furnaces. It is used in medicine as a mild alkali. Magnesium 
hydroxide (Mg(OH) 2 ) is insoluble in water but is a fairly strong 
base. Milk of magnesia is magnesium hydroxide; it is used as a 
laxative. The carbonate (MgC0 3 ) is found as magnesite and with 
calcium carbonate in the mineral dolomite (CaC0 3 .MgC0 3 ) . 
Dolomite is used as a building stone. Magnesium sulfate (MgS0 4 ) 
is found extensively as the hydrate; it is called Epsom salt and is 
used in medicine, in the manufacture of sulfates of sodium and 
potassium, in tanning, for weighting cloth, and in making paints 
and laundry soaps. Serpentine, asbestos, and talc are magnesium 
silicates. Asbestos is used as a covering for pipes, furnaces, and 
boilers because it is a non-conductor of heat. 

Zinc. Like the metals already studied, zinc never occurs free 
in nature. Its principal ores are sphalerite (ZnS) and franklinite 
(ZnO.Fe 2 3 ). Zinc is obtained by roasting the sulfide or carbonate 
to form an oxide (equations, page 445). The oxide is then mixed 
with coal dust and reduced to zinc. The boiling point of zinc is 
low enough to permit the metal to distill over. Commercial zinc 
contains carbon, arsenic, and iron. Pure zinc is made by dis- 
solving this spelter in hydrochloric acid and electrolyzing the zinc 
chloride solution. 

Zinc has a high luster, but tarnishes in the air; at ordinary 
temperatures the action with oxygen is superficial, but at high 
temperatures the metal burns. Water acts with it at high tem- 
peratures. Pure zinc acts slowly with the acids to form hydrogen, 
because the hydrogen collecting on the zinc forms a non-conducting 
coat which protects the metal from coming in contact with the 
electrically charged ions of hydrogen from the acid. If, however, 
impurities are present, like carbon or copper, the hydrogen 



134 ELEMENTARY CHEMISTRY 

molecules are formed on these more negative elements, and the 
zinc is left free to take ionic charges from the hydrogen. For this 
reason copper sulfate, which forms copper, is added in making 
hydrogen with pure zinc. Zinc is used in making alloys, as brass. 
Galvanized iron is sheet iron coated with zinc to protect it from 
rust (page 447); this is the largest use of zinc. Sheets of zinc are 
used for lining sinks. In electric batteries zinc forms one of the 
poles, and it is used in separating silver from lead (page 516). 

Compounds of zinc. The oxide (ZnO), known as zinc white, 
occurs in nature in impure form; it is prepared by burning zinc. 
Zinc oxide finds a large use as a white pigment. It is not colored 
dark by sulfur compounds as white lead is. It is used as a filler 
in rubber goods, such as automobile tires. Zinc sulfate (ZnS0 4 ) 
forms a hydrate known as white vitriol. It is used in electric 
batteries. Zinc chloride (ZnCl 2 ) has an antiseptic action and is 
used to preserve wood from decay, as in railroad ties. Copper 
sulfate and coal-tar creosote are used for this purpose also (page 
448). 

EXERCISES* 

1. Discuss the properties of the alkaline-earth metals. 

2. Give the mineral name, chemical name, and formula for five cal- 
cium minerals. 

3. What is lime? How is it prepared? What are its uses? 

4. What is slaked lime? Give its uses. 

5. What is mortar? Explain how it sets. 

6. How is bleaching powder made? What is its formula? How does 
it act as a bleaching agent? 

7. Name five natural varieties of calcium carbonate. What is pre- 
cipitated chalk? Write the equation for making it. 

8. What is plaster of Paris? How is it made? What are its uses? 

9. What is hard water? Give the cause of the two kinds of hard 
water. 

10. How is hard water softened? 

11. What causes boiler scale? Give its composition. 

12. How is calcium carbide made? What are its uses? 

13. How is calcium cyanamide made? What is its use? 

14. Give the flame coloration imparted by calcium, strontium, and 
barium salts. 

15. What is a fertilizer? Name the three constituents of a fertilizer 
and their chief sources. Is lime a fertilizer? 



*Prepare this set of Exercises and hold it until those for Lessons XXII, XXIII, XXI V, 
and XXV are also prepared and then send all five sets to the School. 



ELEMENTARY CHEMISTRY 135 

16. Give the properties of the magnesium family. 

17. How is magnesium made? What are its uses? 

18. Describe the metallurgy of zinc. 

19. State the properties and uses of zinc. 

20. How is pure zinc made? 

21. Why is zinc used in galvanizing iron? How is iron galvanized? 

22. How is wood preserved? 

23. Give chemical name, formula, and use for each of following: white 
vitriol, sphalerite, zinc white, Epsom salt, gypsum, quicklime, and dolomite. 

24. Answer questions 4, 6, 7, 8, 11, 12, and 13, page 438. 

25. Solve problems 14, 15, 17, 19, and 20, page 439. 

26. Answer questions 5, 6, 7, 8, 9, 11, and 12, page 450. 

27. Solve problems 14, 15, 16, and 17, page 450. 



LESSON XXII 

COLLOIDS; ALUMINIUM 
COLLOIDS 
Assignment: Chapter XXX, McPherson and Henderson 

The colloidal state. Substances which assume a jelly-like 
form are known as colloids. The word means glue. Many 
organic substances are of this nature, such as gelatin, soap, egg- 
yolk, dyes, and the casein of milk. Silicic acid, hydroxides, like 
aluminium, and sulfides, like arsenic and nickel, are inorganic 
substances that tend to assume the colloidal state. 

There are two kinds of colloidal states. One, in which the 
substance seems to be in solution, the other in which it forms a 
jelly with the water. A dilute solution of water glass, to which 
dilute hydrochloric acid is added, remains apparently unchanged 
for a time, but, on standing, a jelly-like precipitate of silicic acid 
forms. The clear liquid just referred to is not a true solution of 
silicic acid. This is shown by passing a ray of light through it in 
a dark room (page 383). The path of the light is visible, as it is 
if passed through the air in a dark but somewhat dusty room. 
The path of the light is visible because small particles of matter 
reflect the light. A true solution, as sodium chloride, shows no 
visible path under similar conditions. 

The first colloidal state, where the substance appears in near 
solution, is called sol, or hydrosol. The second state, in which the 



136 ELEMENTARY CHEMISTRY 

jelly-like form exists, is known as gel, or hydrogel. The essential 
difference between a true solution and a hydrosol seems to be due 
to the size of the particles. In true solutions the particles 
(molecules) are so small as not to be able to reflect the light as 
the rays pass through. In the next state they are large enough 
to reflect some of the light but not large enough to interfere with 
the transparency of the liquid. This is the state of colloidal solu- 
tion. In this state, the liquid does not obey the laws of freezing- 
point lowering as do true solutions. In another state the particles 
are large enough to give an opaque or milky appearance to the 
liquid, but too small to be filtered off or to settle. This may be 
called a colloidal suspension. Suspensions settle after a time, and 
precipitates settle quickly. The difference, then, is a matter of 
the size of the particles. 

Changing conditions can cause the hydrosol to become a 
hydrogel. This change is called coagulation. Some coagulations 
are reversible and some are not. Gelatin and water form a 
hydrosol when warm and a hydrogel when cold as often as desired. 
The white of an egg beaten up with water is a hydrosol. Heating 
coagulates it, and this change is not reversible. The addition of 
an electrolyte will cause the coagulation of some colloids which 
have electric charges upon them. Colloidal gold is a negative 
colloid and is coagulated by the positive ion of a salt; ferric 
hydroxide is a positive colloid and the negative ion of the salt 
coagulates it (page 383). 

Colloids are prepared by two methods: (1) by powdering 
solids; (2) by imperfect precipitation. See the illustration and 
examples on page 384. When hydrogen sulfide is passed into a 
solution of arsenious acid, the sulfide (As 2 S 3 ) is formed, but remains 
in colloidal suspension. It is coagulated by adding an electrolyte 
in the form of a solution of a salt. 

Colloids are not a special kind of substance, but any sub- 
stance may assume the colloidal state, just as it may assume the 
liquid or solid state or the state of a true solution. No sharp 
distinction can be drawn between the different classes. In colloids 
the molecules are bunched together without any order; in crystals 
the molecules assume an orderly arrangement. In the hydrogels 
the bunches of molecules inclose water, forming the jellies. 



ELEMENTARY CHEMISTRY 137 

Many important industries are based upon the properties of 
colloids. Soap solution is a colloid and owes its cleansing action 
mostly to this property. The decolorizing of sugar by charcoal is 
due to the colloidal nature of the coloring matter, the colloid 
being absorbed by the charcoal. Colored glasses have colloidal 
matter suspended in them. Silver salts are colloidally suspended 
in collodion. The use of aluminium hydroxide in the purification 
of water (page 457) and as a mordant for dyes (page 458) depends 
upon its colloidal nature. Rubber is a colloid, as are gums, 
waxes, glues, cements, and bread dough. 

A similar condition exists in emulsions. Here we have two 
liquids not soluble in each other, as oil and water. If the par- 
ticles can be kept small enough, a permanent emulsion results. 
If oil and vinegar are beaten together (French dressing), we have 
a temporary emulsion. If the yolk of an egg (colloidal) is added, 
a permanent emulsion (mayonnaise dressing) results. Milk is a 
somewhat imperfect emulsion as the cream separates out. For the 
emulsion to be stable, it is necessary that a third substance of a 
colloidal nature be added. Soap solution admirably serves the 
purpose of keeping the small particles of oil from joining to form 
larger ones and so keeps oil emulsified. 

ALUMINIUM 
Assignment: Chapters XXXVI and XXXVII, McPherson and 
Henderson 

The group. Aluminium is a member of Group III. Boron is 
the non-metallic member of the group and has already been 
studied. Aluminium is the only metal of the group that occurs 
abundantly, the others being rare, and only aluminium will be 
studied. It is a very useful metal and many of its compounds 
find wide application The metals of the group are trivalent. 
They form weak bases, many of whose salts are hydrolyzed in 
solution. Their oxides are known as earths. 

Occurrence. Aluminium is the most abundant metal, more so 
than iron, but it is difficult to obtain it from some sources. It is 
mostly found in the form of silicates which are hard to decompose. 
Feldspar and clay both contain aluminium, but it is not obtained 
from these sources in any quantity. Bauxite (A1 2 3 .H 2 and 



138 ELEMENTARY CHEMISTRY 

AI2O3.3H2O) and cryolite (Na 3 AlF 6 ) are used in preparing it com- 
mercially. The element is called aluminum as well as aluminium} 
the former word is most used in commerce. 

Preparation. The metal is prepared by the electrolysis of a 
solution of bauxite in fused cryolite. This process is described on 
page 453. The method was devised by Hall in 1886. Prior to 
that time the metal was rare. Cheap aluminium depends upon 
cheap electric energy, and therefore its preparation is one of the 
electrochemical industries at Niagara Falls. 

Properties and uses. Aluminium is a light, white metal. It 
is strong and a good conductor of heat and electricity. Moist air 
tarnishes it slightly, forming a thin film of oxide which prevents 
further action; boiling water has little, if any, effect upon it for 
the same reason. The metal will burn in air. It unites with 
oxygen at high temperatures liberating much heat (page 454). It 
is, therefore, a good reducing agent, and as such it is used in the 
metallurgy of some metals— the Goldschmidt reduction process. 
The thermite process, used for welding steel rails, is an application 
of the same principle. It is described on page 455. 

Dilute acids act upon it, liberating hydrogen, though sulfuric 
acts very slowly and nitric almost not at all, because of the 
insolubility in these two acids of the film of oxide. Concentrated 
and hot sulfuric acid acts with it as with other metals: 

2A1+6H 2 S0 4 ->A1 2 (S0 4 )3+6H 2 0+3S0 2 

Concentrated nitric acid has almost no action on it, but alumin- 
ium in alloys dissolves readily in slightly diluted nitric acid. 
Strong bases react with it to form aluminates and hydrogen: 

2Al+6KOH->2K 3 A10 3 +3H 2 

Aluminium is well adapted to many construction purposes, 
because of its lightness, strength, and non-corrosiveness. It is 
used in the construction of airplanes; cooking vessels are now 
largely constructed from it; as a conductor of electricity it is 
used as trolley wires. A powder form is used as a paint for 
metal surfaces. Alloys are made with it. With copper it forms 
aluminium bronze and with magnesium it forms magnalium. 

An interesting use of aluminium is in cleaning silverware. 
The tarnish on silver is usually silver sulfide. When silver 



ELEMENTARY CHEMISTRY 139 

spoons, knives, etc., are placed in contact with aluminium, as in a 
stew pan, and covered with water, which is then boiled, the 
aluminium replaces the silver, first forming aluminium sulfide. 
This at once hydrolyzes to form aluminium hydroxide. The 
silverware is left with a clean surface, but needs polishing to give 
it the silver luster. 

Compounds of aluminium. Aluminium oxide (A1 2 3 ) occurs 
as emery or in pure form as corundum. Colored with impurities 
it forms precious stones, as sapphire, ruby, topaz, and amethyst. 
Artificial sapphires and rubies have been made which are said to 
be equal to natural stones as jewels. All forms are hard, and 
emery is used for grinding, as in emery wheels. 

Aluminium hydroxide (Al(OH) 3 ) forms as a colloidal precipi- 
tate when ammonium hydroxide is added to an aluminium salt in 
solution. It is insoluble, but hard to filter. Heating drives out 
water, forming the oxide. The special property of aluminium 
hydroxide to be noted is that it will react with both acids and 
bases; like antimony hydroxide, it is amphoteric. See the equa- 
tions on page 457 showing its action with hydrochloric acid and 
with sodium hydroxide. With the latter it forms salts of sodium 
known as aluminates (Na3A10 3 ). 

Aluminium sulfate (A1 2 (S0 4 )3) is cheap and is used in water 
purification, in making alum, in dyeing, and in making paper. 
Alums are double sulfates; two molecules of sulfuric acid have 
three of their hydrogen atoms replaced by aluminium and one by 
potassium to form common alum. Other trivalent elements 
replace aluminium, and other alkali metals replace potassium to 
form other alums (page 459). 

Aluminium nitride (A1N) is formed when the metal and nitro- 
gen are heated together at high temperatures. The nitride reacts 
with water to form aluminium oxide and ammonia (page 461). 
This reaction serves as a method for fixing atmospheric nitrogen. 
Other methods of utilizing atmospheric nitrogen have been men- 
tioned in previous lessons; all of them are summarized on pages 
461 and 462. 

Uses of aluminium hydroxide. Because of its colloidal nature, 
aluminium hydroxide is valuable in purifying water (pages 71 and 
457). The coagulation of the colloid causes it to settle, and as it 



140 ELEMENTARY CHEMISTRY 

does so, it carries with it suspended matter. The hydroxide is 
formed by dissolving in the water some cheap salt which hydro- 
lyzes. Usually enough basic matter is in the water to combine 
with the sulfuric acid set free in hydrolysis: 

Al 2 (S0 4 )3+6H 2 0->2Al(OH)3+3H 2 S0 4 

Otherwise a base, like calcium hydroxide, must be added. 

Aluminium hydroxide is widely used in the dyeing industry. 
Either because of its colloidal nature or because of its amphoteric 
property, it absorbs or combines with many of the dyes (which 
are either weak acids or weak bases). The dyes are organic com- 
pounds prepared from coal-tar derivatives. Many are not fast 
dyes with cotton; in such a case the cloth is dipped in a solution 
of an aluminium salt and exposed to steam. This forms alumin- 
ium hydroxide, which in one way or another holds the dye fast. 
A substance which serves to fix the dye on the fiber is called a 
mordant. 

Baking powders. As was seen in the preceding paragraph, 
aluminium salts hydrolyze. This reaction is complete if a salt of a 
weak acid is used. When a solution of sodium carbonate is added 
to one of aluminium chloride, the expected aluminium carbonate 
does not precipitate. If it forms, it is hydrolyzed to aluminium 
hydroxide (equations, page 460). Use is made of this reaction in 
making alum baking powders. 

Baking powders are of three types. They all contain starch 
or flour and sodium bicarbonate. The soda furnishes the carbon 
dioxide; the starch keeps the powder dry and serves to dilute it. 
They differ in the constituent which is used to liberate the carbon 
dioxide. Alum, cream of tartar, and calcium acid phosphate are 
the three substances used for this purpose. The equations in the 
three cases are as follows: 

2KAl(S04)2+6NaHC03-^2Al(OH)3+3Na 2 S0 4 +K 2 S0 4 +6C0 2 

KHC 4 H 4 6 4-NaHC0 3 ->KNaC 4 H 4 6 +H 2 0+C0 2 
CaH 4 (P0 4 ) 2 +2NaHC0 3 ->CaHP0 4 +Na 2 HP0 4 +2H 2 0+2C0 2 

Baking powders are used to furnish carbon dioxide, which, expand- 
ing, makes the dough light. Sometimes soda and sour milk are 
used. 



ELEMENTARY CHEMISTRY 141 

Aluminium silicates. A number of these silicates, which are 
found in nature and are useful substances, are mentioned on page 
464. Clay products find a wide use in building operations and in 
the manufacture of pottery. Common brick and tile are made 
from clay without chemical change, except the conversion of iron 
compounds to ferric oxide, when the clay is heated. This gives 
these objects a red color. Vitrified brick is made by being heated 
high enough to partially fuse the clay, forming a kind of 
glass. 

Pottery. Pottery includes fine chinaware and crude porce- 
lain. The methods of manufacture are essentially the same and 
consist of three steps: (1) preparation of the body, or bisque; (2) 
glazing; (3) decorating. The different kinds of bisque, from very 
porous to non-porous, are made by varying the quantities of clay 
and feldspar used. The glaze is used to render the ware non- 
porous and to give a smooth surface; it is a fusible glass melted 
into the body of the object. Silica, feldspar, and metallic oxides 
are used in making the glaze. They are made into a paste into 
which the bisque is dipped; the objects are then fired in a kiln to 
fuse the glaze to the body. Metallic oxides or colored glasses 
made from metallic oxides and silica, that is, silicates, are the pig- 
ments used in decorating china. 

Cement. Portland cement is the most important cement. It 
is made by powdering limestone and clay or shale together and 
heating the mixture to a clinker in a furnace. This clinker is then 
ground to a powder and is the cement. Gypsum is sometimes 
added to retard the setting. Cement is a silicate, principally of 
calcium and aluminium. Smaller amounts of iron and magnesium 
silicates are present (table on page 466). Any materials furnish- 
ing these ingredients can be used to make cement. Blast-furnace 
slag is frequently used. Cement will harden under water as well 
as in air. This setting of cement is due to reactions not fully 
understood. Apparently hydrolysis takes place first, and then 
hydrates are formed which crystallize into a hard mass. The use 
of cement is growing greater every year. Mixed with crushed 
stone, it forms concrete, which is used for various kinds of con- 
struction, paved roads, bridges, walls of buildings, and even for 
hulls of ships. 



142 ELEMENTARY CHEMISTRY 

EXERCISES 

1. What is meant by a colloid? 

2. What is a hydrosol; a hydrogel? 

3. What is the nature of the colloidal state? 

4. How are colloids prepared? 

5. Give all the industrial applications of colloids you can. 

6. What are emulsions? Give examples. 

7. Describe the Hall method for making aluminium. 

8. What are the properties and uses of aluminium? 

9. Describe the thermite welding process. 

10. What are alum, corundum, ruby, alundum, sapphire, emery, baux- 
ite, cryolite, feldspar, mica, clay, and Fuller's earth? 

11. Explain the use of aluminium hydroxide in purifying water; as a 
mordant in dyeing. 

12. Discuss the three kinds of baking powders. 

13. Give all the methods for fixing atmospheric nitrogen. 

14. Discuss the manufacture of white pottery. 

15. Discuss the manufacture and setting of cement. 

16. Answer questions 2, 3, 4, 5, 7, and 9, page 462. 

17. Answer questions 10, 11, 12, 13, and 14, page 463. 

18. Solve problems 16 and 17, page 463, and 7, page 468. 

19. Answer questions 1, 2, 3, 4, and 5, page 468. 

LESSON XXIII 
IRON 
Assignment: Chapter XXXVIII, McPherson and Henderson 

The family. Iron, cobalt, and nickel occur in the eighth 
group of the periodic system. Their atomic weights are nearly 
equal and their properties very similar. They have a valence of 
two and three in their compounds. Iron is by far the most 
abundant and important of the three. In fact, it is the most 
widely used of all the metals and is more abundant than any 
other except aluminium. It is mostly used in the metallic form 
as is nickel. These elements are not found free in nature, but in 
the form of oxides, sulfides, and carbonates. Elementary iron has 
been found in meteorites; few rocks or soils are free from iron; it 
is a constituent of the haemoglobin of the blood and of the 
chlorophyll of plants. 

Iron. Pure iron is seldom prepared and is too soft for most 
uses. It can be prepared by the electrolysis of iron sulfate or by 



ELEMENTARY CHEMISTRY 143 

reducing iron compounds with hydrogen. Practically pure iron 
can be obtained by the open-hearth method (page 478). This is 
especially adapted to use in electromagnets. The iron of com- 
merce contains small percentages of other elements, which modify 
its properties somewhat. Carbon, silicon, phosphorus, manganese, 
sulfur, and oxygen are found in iron. Carbon is always present 
in varying amounts and in varying forms: as graphite, which has 
crystallized from the iron on cooling; in solid solution in the iron; 
as carbides, such as Fe 3 C. These carbides may be in solid solu- 
tion or separated out. The state of the carbon depends upon 
the conditions of manufacture and cooling of the iron. For these 
reasons iron exists in three recognized varieties, cast iron, wrought 
iron, and steel. 

The metallurgy of iron. In making iron, large and rapid pro- 
duction is a necessary requirement as well as a product of the 
desired composition. Cast iron is the variety that is always 
made directly from the ore. The materials used are of four kinds, 
the ore, carbon, hot air, and flux. The products of the operation 
are cast iron, slag, and the gas. 

1. Ores. The ores most used are oxides and carbonates of 
iron (page 471). A common iron mineral (FeS 2 ) is used for its 
sulfur in making sulfuric acid, but it is not a satisfactory iron ore, 
because it is difficult to get iron of sufficient purity from it. The 
ores of iron contain impurities which may be classified as acidic 
or basic in nature. The oxides usually have acidic impurities, as 
sand, clay, etc. The carbonate ore contains other carbonates, as 
limestone, and is basic. Sulfides and phosphates are often present. 

2. Carbon. Carbon is used as the fuel and reducing agent. 
However, the carbon does not directly reduce the iron oxide. It 
first burns to carbon dioxide. This is reduced to carbon monoxide, 
which in turn reduces the iron oxide: 

Fe 3 4 +4CO->3Fe+4C0 2 

This carbon dioxide is again reduced by the carbon to the mon- 
oxide. The monoxide mixed with nitrogen escapes from the 
furnace. Coke is the form of carbon usually used, although 
formerly wood charcoal was used. Coal cannot be used as it 
uses up too much heat in giving off its volatile matter. 



144 ELEMENTARY CHEMISTRY 

3. Hot air. Hot air is used to maintain a high temperature 
by supplying oxygen to burn the fuel. A strong blast of hot air 
is forced into the lower part of the furnace during its operation. 

4. Flux. In order that the iron may flow free from the 
earthy matter, it is necessary that these earthy impurities be 
converted into substances which will melt and flow off. This is 
done by making them into a sort of glass. Glasses are silicates 
of metals; that is, they are salts composed of a basic and acidic 
constituent. The nature of the flux needed to form this glass, 
then, depends upon the nature of the impurities in the iron ore. 
If the ore contains sand or clay, it is necessary to add a basic 
flux, such as limestone (CaC0 3 ). If, however, limestone is the 
impurity in the ore, then an acid flux must be used, as sand or 
feldspar. When heated, the impurity and the flux combine and 
form a liquid glass, or slag. 

The slag is essentially a calcium-aluminium silicate, and is 
difficult to fuse. It is produced in that region of the furnace 
where the temperature is about 1200° C. This is the temperature 
at which the cast iron becomes liquid enough to flow freely. If 
the slag were produced higher up in the furnace, it would contain 
some of the iron as iron silicate, and the iron would be lost. The 
slag is lighter than the iron and floats on top, where it is con- 
tinually drawn off through a small opening in the furnace. The 
iron is tapped off at intervals from an opening lower down. 

The gas, consisting mostly of carbon monoxide and nitrogen, 
is carried off by the downcomer pipes. This gas is used for 
heating the hot blast and as a fuel to operate the engines that 
furnish power. The description of the blast furnace (pages 473 
and 474) should be carefully studied. If opportunity ever offers, 
the student should observe a blast furnace in operation. 

Cast iron. The varieties of iron vary in composition, mostly 
with respect to the carbon content, but this is not a good 
guide in distinguishing them from one another. The carbon con- 
tent of one form overlaps that of another. Some steels contain 
more carbon than some cast irons. Generally, however, cast iron 
contains more carbon than the other varieties. The properties of 
the forms of iron depend much upon their method of manufacture. 
We may say that cast iron is the product of the blast furnace. 



ELEMENTARY CHEMISTRY 145 

Frequently it is not cast in molds, but, while still hot, is con- 
verted into steel and then into rails or other forms for ultimate 
use. Sometimes it is cast in molds by machinery or by allowing it 
to run into troughs made in sand (Fig. 187, page 474). These 
castings are called pigs, because the arrangement of the molds 
suggests suckling pigs and the mother hog. 

Cast iron contains from 2 to 5 per cent of carbon, some- 
times nearly as much silicon, and smaller amounts of manganese, 
phosphorus, and sulfur. Two kinds of cast iron occur, depending 
upon the form in which the carbon exists. If suddenly cooled, the 
carbon, as iron carbide, remains in solid solution. This iron is 
homogeneous and very brittle. It is called white, or chilled, cast 
iron. If the cooling is slower, the carbon separates as graphite, 
which appears as black scales. This form is called gray cast iron. 
Cast iron has a lower melting point than pure iron. It expands in 
casting and is used for making stoves, radiators, and some parts 
for machinery. It is rigid, but not elastic, so breaks easily. It is 
the starting point for making wrought iron and steel. 

Wrought iron. This is the purest of the three varieties. It is 
made by burning out the carbon and other impurities in a pud- 
dling furnace, so called because the iron is stirred with iron rods 
(as in a puddle). As it becomes pure, the iron stiffens and is 
withdrawn in balls on the ends of the stirring rods. Oxide of iron 
is added to the lining of the furnace to supply oxygen. Wrought 
iron melts at a higher temperature than cast iron, is soft, and has 
a fibrous structure, while cast iron has a granular structure. It is 
not produced on a large scale, as steel serves for most of the pur- 
poses for which it could be used. 

Steel. This variety of iron is almost free from silicon, phos- 
phorus, and sulfur, but contains carbon from traces up to 2 per 
cent, as may be desired for different uses. Steel has a very fine 
granular structure. It is elastic and hard; it can be forged, cast, and 
rolled. It is the product of the Bessemer or the open-hearth process. 

The Bessemer process was invented by Kelly, an American, 
and by Bessemer, an Englishman, at about the same time. It is 
usually used at the present time for irons which contain little 
phosphorus; the converter is then lined with fire clay or silica 
(study description on page 477). A modification of this process 



146 ELEMENTARY CHEMISTRY 

has been used for irons containing considerable phosphorus; the 
converter is then lined with lime and magnesia, which will form a 
slag containing the basic phosphate of calcium. This is known as 
the basic-lining, or Thomas-Gilchrist, process. 

The open-hearth process can be used for both kinds of iron. 
In this country the furnace is lined with a basic substance like 
lime and magnesia and is sometimes called the basic open-hearth. 
The furnace can be lined with sand and used for irons that con- 
tain little or no phosphorus. Gas or oil is used as the fuel to 
heat the open-hearth furnace. The basic open-hearth method re- 
moves all but traces of phosphorus and sulfur from the steel. A 
better grade of steel can be made by this process as better regula- 
tion is possible (pages 478 and 479). 

Crucible steel is used in making sharp tools. It is made from 
other steel by melting it in a crucible with the desired amount of 
pure carbon. A high-grade tool steel is made now in electric 
furnaces, the current being used only to melt the steel. 

Tempering. Steel contains carbon in the form of the element 
or in the form of iron carbide (Fe 3 C) dissolved in the iron. When 
the steel is suddenly cooled, no changes take place, and a solid 
solution results. This is hard and brittle. If the steel is cooled 
slowly, some of the carbide (Fe 3 C) separates out as crystals, until 
at 700° C. only 0.9 per cent carbon remains in solution. Under 
this condition, the solid solution forms a mixture of pure iron and 
iron carbide. The pure iron is soft, the carbide is hard. If such 
steel is heated for tempering, the amount of softer material formed 
depends upon the temperature and the rate and the time of cool- 
ing. Steels are tempered to different degrees of hardness for 
different uses (page 481). 

Steel alloys. Many elements as well as carbon are added to 
steel to vary its properties. The elements most frequently used 
are manganese, silicon, nickel, chromium, tungsten, vanadium, and 
titanium. These mixtures are called alloy steels. A list of them 
and their composition and uses is found on page 481. Aluminium, 
vanadium, and titanium are used as steel purifiers because they 
combine with the oxygen and the gases left in the iron, completing 
the reduction and preventing the presence of gases, which causes 
blowholes. 



ELEMENTARY CHEMISTRY 147 

Chemical properties of iron. Iron is above hydrogen in the 
electrochemical series of the metals. It acts readily with acids, 
forming hydrogen. With dilute acids, the hydrogen is almost 
pure; with concentrated acids, the carbide present is affected and 
hydrocarbons are formed. Pure iron does not rust in cold water, 
but ordinary iron does. The rust consists mainly of hydrated 
ferric oxide (3Fe 2 3 .H 2 0) . The presence of carbon dioxide hastens 
the speed of the action. Iron rust is porous and also tends to 
scale off, so it does not protect the metal from further action as 
does zinc rust. Iron burns with oxygen and acts with super- 
heated steam to form Fe 3 4 . A superficial layer of this will pro- 
tect the iron from the action of the air. Duriron and tantiron, 
which are used for acid containers and do not rust, are silicon 
alloys. 

Compounds of iron. Iron forms two series of compounds, one 
with a valence of two, and one with a valence of three. The 
bivalent compounds are called ferrous, the trivalent, ferric. The 
more important iron compounds are described on pages 482 to 
485. An important property of iron compounds is their ability to 
be oxidized from the ferrous to the ferric state and reduced from 
the ferric to the ferrous condition. 

Moist air will change ferrous compounds to ferric: 

2Fe(OH) 2 +0+H 2 0->2Fe(OH) 3 

Any good oxidizing agent may be used as the source of the oxygen, 
such as nitric acid or potassium permanganate. In order to get a 
simple oxidation, the presence of the acid having a common anion 
with the salt is necessary. If nitric acid oxidizes ferrous chlorides 
the equations are: 

2HN0 3 ^H 2 0+2NO+3[0] 
6FeCl 2 +3[0]+6HCl->6FeCl 3 +3H 2 
adding gives 

2HN0 3 +6FeCl 2 +6HCl->4H 2 0+2NO+6FeCl 3 

Chlorine will convert ferrous salts to ferric: 
2FeCl 2 +Cl 2 ->2FeCl 3 
This is termed oxidation just as if oxygen itself were used. The 



148 ELEMENTARY CHEMISTRY 

term oxidation is applied to any reaction when the valence of the 
positive ion is increased. 

Nascent hydrogen, hydrogen sulfide, and some other sub- 
stances will reduce ferric salts to the ferrous state: 

FeCl 3 +[H]->FeCl 2 +HCl 

When the valence of the positive ion is decreased, the action is called 
reduction. Oxalic acid is a reducing agent and removes iron rust 
(ferric oxide) by reducing it to ferrous compounds, which wash 
out. Iron compounds are used in blue printing (page 488) and in 
some inks (page 489). Ferrous sulfate is used as a disinfectant, to 
destroy weeds, and in purifying water. 



EXERCISES 

1. Name the elements of the iron family. What are the family prop- 
erties? 

2. How is pure iron made? What are its properties? 

3. Discuss the occurrence of iron and name the principal iron ores. 

4. Discuss the four materials used in making cast iron. 

5. How is the kind of flux to be used determined? 

6. Describe the blast furnace. 

7. What is the slag? What is its use in the blast furnace? 

8. What purpose does carbon serve? What chemical changes take 
place in reducing the iron? 

9. What gas escapes from the blast furnace? What is its use? 

10. State the properties of cast iron. How are the two kinds formed? 

11. What are the properties of wrought iron? How is it made? 

12. Describe the Bessemer process for making steel. 

13. Describe the open-hearth process. What are its advantages? 

14. How is crucible steel made? What is its use? 

15. Discuss the tempering of steel. 

16. Discuss the steel alloys. 

17. Give the chemical properties of iron. 

18. Give two examples to show the oxidation of ferrous to ferric iron. 

19. Define oxidation and reduction in this broader sense. 

20. What are the properties and uses of nickel? Describe nickel- 
plating. 

21. Answer questions 2, 3, 5, 6, 8, and 11, pages 491 and 492. 

22. Solve problems 4, 12, and 13, pages 491 and 492. 



ELEMENTARY CHEMISTRY 149 

LESSON XXIV 
THE COPPER AND PLATINUM FAMILIES 
Assignment: Chapters XXXIX and XLIII, McPherson and 
Henderson. 

Introductory. In this lesson we will study the metals copper, 
silver, and gold, which belong to the second family of Periodic 
Group I. Mercury will be studied with these metals, because, 
owing to its position in the electrochemical series, it is very similar 
to them in properties, though not of their family. Gold, while 
belonging to this family, is usually studied with the platinum 
metals, which find a place in Group VIII. Gold is like platinum 
in its action toward oxygen and the acids. 

All these metals may be properly considered together because 
of their non-activity toward dilute acids. They are all suffi- 
ciently inactive so that they are found in the free state in nature. 
Copper, silver, and mercury are dissolved by the oxidizing action 
of nitric acid and hot concentrated sulfuric acid, while gold and 
platinum metals are not. They are dissolved by the action of 
aqua regia, however. Silver is a univalent metal. Copper and 
mercury form univalent and bivalent compounds; gold acts as a 
univalent and trivalent element; platinum is bivalent and quadri- 
valent. Silver, gold, and platinum are known as the "noble metals." 
Metallurgy of copper. Copper ores are of three kinds: (1) 
native copper, which occurs in the Lake Superior region; (2) 
oxygen ores, which are the oxides and carbonates, as cuprite 
(Cu 2 0) and malachite (CuC0 3 .Cu(OH) 2 ) ; (3) sulfur ores, which are 
frequently double sulfides like bornite (Cu 3 FeS 3 ) but are not 
always such since chalcocite (Cu 2 S) is found also. 

Native copper ore is ground and washed to separate the rock 
from the copper. The powdered copper is then melted with a 
flux. The carbonates form oxides when heated and the oxides 
must be reduced with coal. It is more difficult to obtain copper 
from sulfur ores, as they cannot be reduced by heating with 
carbon. Sulfur ores must be heated with a siliceous flux to con- 
vert the iron to a slag. The matte, a mixture of copper and iron 
sulfides, collects under the slag. Sometimes the matte is made in 
a small blast furnace. The matte is converted to blister copper 



150 ELEMENTARY CHEMISTRY 

by a process similar to the Bessemer process for iron. The iron 
forms a slag with the silica, which is added. Blister copper must 
be refined by electrolysis (page 496). 

Properties and uses of copper. The uses of copper depend 
upon its inactivity toward water, dilute acids (except nitric), and 
oxygen at ordinary temperatures and upon the fact that it is 
malleable, ductile, and a good conductor of electricity (page 497). 
Copper is an important constituent of many alloys (see table, page 
497). In addition to the uses mentioned on page 497, it was 
formerly used for making tools and castings; it is used sometimes 
for making kettles, boilers, and tubes for washing machines; great 
quantities are used in electrical plants and appliances; it is used in 
electrotyping. The United States furnishes over half of the copper 
used. 

Compounds of copper. Cuprous and cupric compounds are 
known. Copper is univalent in the former and bivalent in the 
latter. The cupric compounds are much the more common and 
important. Cupric sulfate, when crystallized with water of hydra- 
tion (CuS0 4 .5H 2 0) is called blue vitriol. It is used in copper- 
plating; in batteries; as a mordant in dyeing; to destroy algae in 
drinking water; mixed with milk of lime (Bordeaux mixture), to 
spray grape vines, fruit trees and vegetables; in making insecti- 
cides; in copper refining; and as a remedy in hoof disease. The 
use of copper sulfate in electric cells (page 500) should be given 
special attention. 

Mercury. The chief ore of mercury is cinnabar (HgS). It 
occurs native as well. It is easily obtained by heating the ore 
with carbon; the mercury is volatile and boils over. It is a liquid 
at ordinary temperatures. It acts much like copper toward acids 
and water but is somewhat less active toward oxygen. The metal 
is used in thermometers, barometers, and in making alloys; its 
alloys are called amalgams. It is used in the metallurgy of gold 
and silver, extracting them from the crushed ores by forming 
amalgams. The finely divided element is used in medicine (blue pills) . 

Two series of salts are formed, mercurous (univalent) and 
mercuric (bivalent). The chlorides are both important. Mercu- 
rous chloride is the common drug calomel (HgCl). It is a white 
insoluble salt and is used as a purgative, usually mixed with 



ELEMENTARY CHEMISTRY 151 

sodium bicarbonate to prevent acids from converting it to the 
mercuric form. Mercuric chloride (HgCl 2 ), corrosive sublimate, is 
very poisonous. It is used in dilute solutions as an antiseptic in 
dressing wounds and as a means of killing bacteria in labora- 
tories. Serious cases of poisoning occur by taking the salt by 
mistake. Mercuric sulfide (HgS) occurs in a black and a red 
form. The red form, known as vermillion, is used as a pigment 
in paints. Mercuric fulminate (Hg(ONC) 2 ) is used as the explo- 
sive in percussion caps. 

Silver. Native silver is found alloyed with gold and copper, 
usually scattered through rocky material. Silver sulfide (Ag 2 S) is 
the chief source of silver. It occurs alone or mixed with lead 
sulfide (PbS). The metallurgy of silver is dependent upon the 
substances with which it occurs. If alloyed with gold, it is a 
product of the extraction of gold (page 550) ; if it occurs as a 
sulfide with lead, it is obtained in crude form from the lead by the 
Parkes process (page 516). The crude silver is refined by parting 
with sulfuric acid or by electrolysis. If sulfuric acid is to be 
used, the silver is first cupelled (page 504) to oxidize the baser 
metals; sulfuric acid dissolves the silver from the gold; finally the 
silver is precipitated from the sulfate by copper. In refining by 
electrolysis silver nitrate is used as the electrolyte. 

Properties of silver. While silver is similar to copper in 
many respects, it does not act with oxygen. Ozone will oxidize it. 
It is below copper in the electrochemical series and generally is 
less active, yet its oxide is quite basic, absorbing carbon dioxide 
from the air. Its oxide, like that of mercury, can be decomposed 
by heat. Sulfur and sulfur compounds (eggs, mustard, rubber, 
perspiration) quickly tarnish silver. Dilute acids and fused alka- 
lies do not act with it. It is used in alloys for coins and table- 
ware, and mirrors are backed with it. Many objects made of 
cheaper metals are plated with silver by electrolysis; potassium 
silver cyanide is usually used as the electrolyte. 

Silver compounds. The nitrate of silver (AgN0 3 ) is known 
as lunar caustic. It is used in surgery because of its corrosive 
action on flesh. The halogen compounds with silver are used in 
photography. The chloride, bromide, and iodide of silver are all 
insoluble in water and undergo a change of color and composition 



152 ELEMENTARY 7CHEMISTRY 

when exposed to sunlight. The bromide and iodide are most used 
and require a developer (reducing agent) to bring out the image. The 
developers are organic compounds. (See "Photography/' page 508.) 

Gold. Most of the gold is found in the free state. There are 
three ways of mining gold. Placer (c is pronounced like s) 
mining is the oldest method, but hydraulic and vein mining are 
more used now (page 550). The extraction of gold from its ore 
can be accomplished by the amalgamation or the cyanide proc- 
ess (page 550). The cyanide process is successful with ore con- 
taining small percentages of gold. Sometimes it is used to extract 
gold from the "tailings" left from other processes. The introduc- 
tion of the cyanide process greatly increased the gold production 
of the world. Gold extracted in these ways contains other metals, 
as silver, lead, copper. It is refined by electrolysis, cupellation, 
parting with sulfuric acid, or by a combination of the last two 
methods (pages 550 and 551). 

Properties of gold. Gold is an inactive element. It does not 
combine with oxygen, nor the common acids. Aqua regia dis- 
solves it, and fused alkalies attack it. It reacts with free chlorine 
and bromine. It is a soft metal, and to give it hardness, it is 
alloyed with silver or copper. Gold compounds are decomposed 
by heat. Other metals replace gold from solutions of its salts, and 
reducing agents likewise precipitate the gold. It is precipitated in 
the form of a brown powder. Mixed with oil, it is used in this 
form for decorating china. When fired, the gold assumes its 
common yellow color. It is used mostly for coins and jewelry. 

The platinum family. Two sets of three metals each occur in 
Group VIII after the iron family. The first have atomic weights 
about 100. Palladium is the most important of the three. It is 
distinguished by its ability to absorb large quantities of hydrogen 
and it is used for this purpose in gas analysis. The next set have 
atomic weights near 200. Platinum is the m@st important of this 
set. All six are found alloyed together in nature and closely 
resemble each other in properties. They are known as the plati- 
num family. 

Platinum. Most of our platinum comes from the Ural 
mountains in Russia. It is separated from the sand by washing. 
Native platinum is 60-84 per cent platinum. The other mem- 



ELEMENTARY _ CHEMISTRY 153 

bers of the family~and gold constitute the rest. To separate the 
platinum, the alloy|.\is dissolved in aqua regia. Ammonium 
chloride is used to precipitate the platinum (page 545). Platinum 
is an inactive element, being acted upon only by aqua regia, the 
fused alkalies, and chlorine, and it has a high melting point. For 
these reasons it is much in demand for making laboratory utensils. 
Large platinum dishes are used to concentrate chamber sulfuric acid. 
It is used as a catalytic agent, and is prepared in a very finely 
divided condition for this purpose. It is so used in making sul- 
furic acid by the contact process and in making formaldehyde 
from methyl alcohol. Platinum has the same coefficient of ex- 
pansion as glass and was formerly used in light bulbs. It is used 
in jewelry, photography, and dentistry. In fountain-pen points 
an alloy with iridium is used because of its greater hardness. 
Many substitutes are being devised for platinum because of its great 
cost. It is more than twice as costly as gold, and the price is increasing. 

EXERCISES* 

1. Discuss the metallurgy of copper. 

2. State the properties of copper. 

3. How does copper react with hydrochloric acid; dilute sulfuric acid; 
dilute nitric acid; fused alkalies; oxygen; water; hot concentrated sulfuric 
acid? Write equations where action occurs. 

4. Give the uses of copper. Name its alloys. 

5. What are the uses of copper sulfate? 

6. Explain the action of the Daniell cell. 

7. How is mercury obtained from its ore? 

8. Give the properties of mercury; also the uses. 

9. How does it act with the reagents named in question 3? 

10. What are calomel and corrosive sublimate? Give their properties 
and uses. 

11. Discuss the occurrence and the metallurgy of silver. 

12. Give the properties of silver; also the uses. 

13. Give the reactions of silver with the reagents in question 3. 

14. Discuss fully the subject of photography. 

15. Discuss the mining, extraction, and refining of gold. 

16. Give the properties and uses of gold. 

17. How does gold act with reagents in question 3? 

18. Give the occurrence and the extraction of platinum. 

19. Discuss the properties and the uses of platinum. 

20. Give a use for iridium, osmium, palladium, and lunar caustic. 

21. Answer questions 1, 2, 3, 4, and 5, page 509. 

22. Answer questions 6, 7, 8, 9, 10, and 12, page 510. 

23. Solve problems 16, 17, and 18, page 510. 



154 ELEMENTARY CHEMISTRY 

LESSON XXV 

OTHER METALS 

Assignment: Chapters XL, XLI, XLII, and XLIV, McPherson 
and Henderson. 

Introductory. Several metals remain to be studied, and cer- 
tain others need a brief reference. They are discussed in this 
lesson, but they do not belong together in the periodic grouping, 
nor are they closely related in properties. Tin and lead are some- 
what alike, so are manganese and chromium, while radium and 
uranium have a similarity of properties. 

Tin. The principal ore of tin is the oxide, cassiterite (Sn0 2 ). 
Tin does not occur in the free state. Simple reduction with car- 
bon produces the metal. Tin has long been known, specimens 
being found in the Egyptian tombs. 

It occurs close to, but above, hydrogen in the electro- 
chemical series. It is not so active as zinc or iron, but is more 
active than copper or silver. Dilute acids act slowly with it. 
Concentrated hydrochloric acid forms stannous chloride and hydro- 
gen, concentrated sulfuric acid oxidizes it to stannous sulfate (page 
512), and nitric acid forms metastannic acid (H 2 Sn0 3 ). Tin burns 
at high temperatures, but is unchanged by air or water under 
ordinary conditions. It is attacked by fused alkalies. It forms 
bivalent and quadrivalent compounds. Both oxides and their 
hydroxides have acidic and basic properties. Stannous hydroxide 
is mainly a base, and stannic hydroxide is mainly an acid. The 
salts of stannic acid are called st annates. 

Because of its inactivity, tin is used to plate sheet iron for 
roofing and tinware. Tin does not rust, but if scratched off, the 
iron rusts faster than if no tin were in contact with it (page 512). 
With galvanized iron, the zinc is attacked first. Tin is used in 
many alloys (page 497). Tin is recovered from tin cans and tin 
scrap by detinning with chlorine (page 514). 

Lead. The principal ore of lead is the sulfide, galena, or 
galenite (PbS). Sulfide ores must be roasted before they can be 
reduced. The process can be so conducted that lead sulfide 
serves to reduce the oxides and sulfates formed by the roasting 



ELEMENTARY CHEMISTRY 155 

(equations, page 515). Silver is obtained in the process, alloyed 
with the lead. Hard lead is first obtained, the hardness being 
due to impurities of arsenic, bismuth, antimony, copper, etc. 
These are burned to oxides, which float, by melting in an open 
furnace. The soft lead, thus obtained, contains silver. This is 
removed by the Parkes process (page 516). 

Lead is a heavy metal, but not so heavy as gold, platinum, 
or mercury. It is just above hydrogen in the electrochemical 
series. It tarnishes, but is not acted upon to any extent by 
oxygen of the air. At high temperatures it burns. Nitric acid 
oxidizes it, but hydrochloric acid and sulfuric acid have little 
action on it because of the formation of insoluble lead salts. 
Weaker acids, like acetic, act with it. Lead is used in many 
alloys, in storage batteries, in pipes for plumbing, and in making 
white lead. Litharge (PbO), red lead (Pb 3 4 ), and the peroxide 
(Pb0 2 ) are oxides of lead. Their uses are given on page 518. 

Paints. Paints consist of three ingredients, body, vehicle, and 
pigment. White lead, basic lead carbonate (2PbC0 3 .Pb(OH) 2 ) is 
the most commonly used body material. Zinc oxide is good. 
Barium sulfate and china clay are inferior substitutes, if sold as 
white lead. White lead is made from lead plates, acetic acid, 
and fermenting organic matter (page 519). The best vehicle is 
linseed oil. It drys by absorbing oxygen, and to make it dry 
quicker, it is boiled with oxides, which start the oxidation. The 
pigment is the coloring matter. This may be the body also, as 
lead chromate (yellow), but usually the pigment is a metallic 
oxide. Organic dyes are used in the form of lakes (pages 520 
and 521). 

Storage Batteries. There are two kinds of cells for storing 
electrical energy as chemical energy which can be liberated as 
electrical energy. The chemical action occurring when the cell 
is being charged is reversed when the cell is discharging current. 
The ordinary cell consists of lead plates covered with spongy 
lead at one pole and lead dioxide at the other when ready for 
use. The liquid is sulfuric acid solution. As it discharges, both 
plates become coated with lead sulfate (page 522). The other 
type (Edison cell) has one plate of iron and the other of nickelic 
oxide (Ni 2 03). The solution is potassium hydroxide. The cell 



156 ELEMENTARY CHEMISTRY 

operates by the "nickel oxide being reduced to nickel hydroxide 
(Ni(OH) 2 ) and the_ironJbeing oxidized to Fe(OH) 2 , which action 
delivers energy. 

Manganese. This element forms a great variety of com- 
pounds. It has a valence from two to seven as shown by the 
oxides (page 525). ? Manganese dioxide, pyrolusite (Mnft), is the 
principal source of manganese. It is a good oxidizing agent. 
Manganese acts both as a positive and a negative element. As 
a positive element its manganous salts (page 525) are most 
important. Potassium permanganate (KMn0 4 ) is the most com- 
mon example of its acid-forming property. This compound is 
frequently used as an oxidizing agent (page 527) and is also an 
antiseptic and disinfectant. 

Chromium. Chromium does not belong in the same group 
with manganese, but it acts as an acid-forming and base-forming 
element. It is hard to reduce from its ores, as is manganese. 
Chromium is trivalent and hexavalent. In its hexavalent form 
its acid properties are more noticeable. Chromic acid (H 2 Cr0 4 ) 
and dichromic acid (H 2 Cr 2 7 ) are unstable but form several well- 
known salts. Lead chromate (PbCr0 4 ) is a yellow pigment; 
potassium dichromate (K 2 Cr 2 7 ) is a common oxidizing agent 
(page 532). 

Radioactivity. The term "radioactivity" is applied to the 
power which certain elements have of throwing off rays of a 
peculiar nature. These rays will pass through some substances 
opaque to light rays and affect a photographic plate. Also, they 
discharge a charged electroscope (page 536). Several elements 
and their compounds exhibit this property, uranium, radium, and 
thorium being among them. Radium shows the most marked 
radioactivity. This property was first discovered in connection 
with uranium compounds by Becquerel. The Curies observed 
that pitchblende, a mineral of uranium, was four times as active 
as uranium. This led them to believe that a more active radio- 
active substance existed as an impurity in pitchblende, and they 
succeeded in isolating radium from large quantities of pitchblende 
residues. 

Radium. The element has chemical properties similar to those 
of barium, and it is placed after barium in that family. Radium 



ELEMENTARY CHEMISTRY 157 

disintegrates, even though it is a ivell-defined element. In this dis- 
integration, it forms helium and niton, both of which belong to 
Group 0. Niton further decomposes, and it is thought that the 
final product is lead (page 542). 

In these decompositions two kinds of particles are thrown off. 
The first kind, called alpha rays, consists of helium atoms charged 
positively; the second kind, called beta rays, consists of electrons, 
which are about rwo as heavy as the hydrogen atom and are 
charged negatively. A third kind of radiation consists of waves 
in the ether and not of particles of matter. These are the 
gamma rays (page 538). Note the demonstration of these three 
kinds of rays in Fig. 208, page 539. 

It is thought that radium is being formed from uranium. A 
great deal of energy is given off in the decomposition of radium. 
The rays from radium compounds disintegrate glass and water, 
produce severe burns upon the skin, and kill bacteria; they are 
used in treating some diseases, as cancer, and it is claimed that 
certain forms of cancer, if treated in the early states, can be 
cured. 

The radium atom. The decomposition of radium suggests 
that the atoms of the elements must have a complex structure. 
It seems probable that the atoms are made up of a positive nu- 
cleus consisting of charged atoms of helium (possibly hydrogen 
also) closely packed and making up nearly all the mass of the 
atom and of negative electrons revolving about the nucleus. For 
radium this system is unstable, but for most of the elements the 
atom is stable. Gradually the atoms of radium are breaking up, 
the explosion resulting in vibrations that cause the gamma ray. 

Rare elements. Some of the rare elements have an impor- 
tant use or two worth mentioning here. Thus thorium and cerium 
oxides are used in making gas mantles. Cerium compounds are 
used in photography and in mordants. Cerium oxalate is used in 
seasickness. An alloy of cerium and iron is used as a gas lighter. 
Titanium, vanadium, molybdenum, and tungsten are used in steel 
alloys. Tungsten is used for the filament in electric-light bulbs 
and is replacing platinum for electrical contacts. Vanadium com- 
pounds are used as catalytic agents, as photographic developers, as 
mordants, and in coloring glass. Titanium oxide is used to color 



158 ELEMENTARY CHEMISTRY 

glass yellow and in electric-arc carbons to give a more efficient 
light. Selenium is a non-conductor in the light. Therefore, it has 
been used in automatic fire alarms. It imparts a red color to 
glass and enamels. 

EXERCISES 

1. State the source and give the metallurgy of tin. 

2. Discuss the properties and the uses of tin. 

3. Describe the metallurgy of lead. 

4. What is the Parkes process? 

5. Give the properties and the uses of lead. 

6. Discuss the composition of paints. 

7. Describe the manufacture of white lead. 

8. Give names, formulas, and uses for the oxides of lead. 

9. Explain the action of storage cells. 

10. What is pyrolusite? Wliat are its uses? 

11. Explain the oxidizing action of potassium permanganate. 

12. What is chrome yellow? What is its use? Write the equation for 
its formation. 

13. Explain the oxidizing action of potassium dichromate. 

14. What is radioactivity? Name three radioactive elements. 

15. How was radium discovered? 

16. What kinds of rays are emitted by radium? 

17. What are the uses of radium? 

18. Discuss the structure of the radium atom. 

19. Answer questions 1, 4, 5, 7, 12, and 13, page 523. 

20. Solve problems 8, 9, 11, and 14, page 523. 

SEND EXERCISES FOR LESSONS XXNXXV TO THE SCHOOL 



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